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The universe is at least 13.7 billion years old and more than 100 billion light years in diameter.
We live on Earth which is part of a solar system with a star called the Sun. Earth and its star coalesced about 4.5 billion years ago. Earth’s solar system is part of the Milky Way galaxy, which is at least 12 billion years old, contains over 100 billion stars and 100 billion planets, and is over 100,000 light years in diameter (or about 0.0001% of the size of the universe).
There are over 100 billion galaxies in the universe.
This page assumes basic knowledge of math, algebra, geometry, and scientific notation.
The universe is made of matter which is something that has mass (m) and takes up a position in space.
Examples of matter are protons, neutrons, and electrons. Protons and neutrons each have a mass of about 1.6×10−27 kilograms (kg), whereas electrons have a much smaller mass of about 9.1×10−31 kg. Protons and neutrons each have a radius of about 0.8×10−15 meters (m), whereas electrons have a radius of less than 1.0×10−22 m.
An object (or entity) is a collection of matter. An object has an inertia, meaning:
- An object at rest will stay at rest unless a force (F) is applied to push or pull it, and
- An object in motion will stay in motion unless a force is applied to push or pull it.
The speed of an object is a description of its movement and is the rate at which its position in space changes over a distance over a period of time (e.g. moving at 13 meters per second (s), or 13 m/s).
The velocity (v) of an object is a description of both its speed and its direction in space (e.g. moving at 13 m/s down).
The acceleration (a) of an object is the rate at which its velocity changes over a period of time (e.g. accelerating at 13 m/s over 1 second down, or 13 m/s2 down).
The momentum (p) of an object, or its “quantity of motion”, is its mass multiplied by its velocity (
If the mass of an object is constant, the force acting on it is the rate of change of its momentum over time (
F=dp/dt); or, substituting momentum for acceleration and simplifying, its mass multiplied by its acceleration (
F=m×a). The unit of force is the newton (N).
When objects touch each other, the perpendicular force their surfaces apply is called the normal force (Fn), and any parallel forces are called frictional forces (Ff).
When an object A exerts a force on another object B, then object B exerts an equal and opposite force on object A.
Electric charge is the property of some forms of matter to create an electromagnetic field which applies a positive or negative force on other electrically charged matter, proportional to the distance between them. The magnitude of this force (named the electrostatic force) is calculated with Coulomb’s law (
F=ke×((q1×q2)/r2)) which is Coulomb’s constant (ke) multiplied by the magnitudes of the two charges involved (q1 and q2) and divided by the square of the distance between the two objects (r2). This equation means that positive charges repel other positive charges, negative charges repel other negative charges, and positive charges attract negative charges (and vice versa).
Examples of electrically charged matter are protons, which are positively charged, and electrons, which are negatively charged. Protons and electrons have the same magnitude elementary electric charge, denoted 1e and -1e, respectively. Neutrons have no charge (electrically neutral).
An electromagnetic field creates electromagnetic radiation which is a wave of its force traveling (radiating) through space. A wave represents something that repeats over time at a frequency which is how often it repeats a cycle per unit time (e.g. one cycle per second). If the unit of time is one second, the unit of frequency is Hertz (Hz). Wavelength is the distance a wave covers over one cycle.
The force carrier for electromagnetic radiation is a photon. Photons are considered particles (which may also refer to objects) even though photons are massless. In a vacuum (space without matter), photons travel at the maximum speed of light (c), or approximately 3×108 m/s. Light is simply electromagnetic radiation; although, colloquially, light usually refers to visible light which is the subset of light that humans see. Fluorescent substances absorb ultraviolet radiation and emit visible light.
Energy (E) is the amount of work one object performs on, or transfers to, another object. Energy is often measured in Joules (J). One joule may be defined as
J=N×m or a force of one Newton acting on an object in the direction of its motion for one meter.
Energy is either kinetic energy if an object is in motion (in most simple cases expressed as
(m×v2)/2); or, potential energy, which may be thought of as stored energy.
Energy may neither be created nor destroyed, but only transformed. Mass may be converted to energy and vice versa through the mass-energy equivalence equation (
Pressure and Temperature
Pressure (P) is the normal force per unit area (A) applied to an object or
P=Fn/A, often measured in Pascals (Pa) (in N/m2). An object under pressure has potential energy.
Atmospheric pressure is the pressure of a planet’s atmosphere (e.g. the pressure of air) on an object. The mass of air in the Earth’s atmosphere decreases expontentially with altitude so the atmospheric pressure decreases expontentially as an object rises above sea level.
Another unit of pressure is the atmosphere (atm) which is the pressure on Earth at sea level and it’s equivalent to 101,325 Pa.
Heat is a measure of the total quantity of kinetic energy. Temperature is a measure of the average kinetic energy of a set of objects. Absolute zero is the coldest state at which an object has minimal movement; however, practically, it may be assumed that all objects have some movement (or vibration) and thus some non-zero temperature.
Temperature or heat is either measured in degrees (°) of change on the scale of Celsius (°C) or Fahrenheit (°F), or in absolute terms on the scale of Kelvin (K).
Roughly, the scale of celsius is defined with 0°C being when water freezes at 1 atm, 100°C when water boils at 1 atm, and absolute zero is -273.15°C. The scale of fahrenheit is defined with 32°F being when water freezes at 1 atm, 212°F when water boils at 1 atm, and absolute zero is -459.67°F. Originally, the Kelvin scale was defined relative to Celsius, but is now defined as pressure invariant, with 0 as absolute zero, and 273.16K when water reaches its triple point. Kelvin was designed so that an increase of one Kelvin is equal to an increase of 1°C.
The lowest recorded surface temperature on Earth is
184K / -89.2°C / -128.6°F, the highest
331K / 58°C / 136.4°F, and the average
288K / 15°C / 59°F.
The specific heat is the amount of heat that must be absorbed or lost for 1g to change its temperature by 1°C.
All matter above absolute zero temperature continuously emits some of its kinetic energy as photons of electromagnetic radiation called thermal radiation or heat. The temperature determines the emission spectrum of wavelengths of the electromagnetic radiation. Most matter may also absorb some of any incoming electromagnetic radiation.
If two objects touch with a path permeable to heat, then, all else being equal, the hotter object heats the cooler object through thermal conduction (and thermal convection in the case of fluids and gases) until they (or at least their touching surfaces) reach thermal equilibrium.
Atoms are made of at least one proton and zero or more neutrons in the atom’s nucleus (both protons and neutrons may be called nucleons), and zero or more electrons orbiting the nucleus in electron shells. Electron shells are also called energy levels because the shell number represents the relative potential energy of electrons in that shell. As electrons are in shells farther from the nuclear proton(s), their potential energies increase because of the additional potential electrostatic force. Therefore, an electron moving to a shell closer to the nucleus must lose energy (e.g. heat) and an electron moving to a shell farther away must absorb energy (e.g. light). Shells are subdivided into subshells, and subshells are subdivided into orbitals (discussed later).
The number of nucleons is considered the atomic mass number (A) since an atom’s electrons’ masses are so much relatively lighter than the nucleons’ masses.
The number of protons is considered the atomic number (Z) and categorizes the atom in a class called a chemical element (e.g. Carbon is the chemical element class for any atom which has 6 protons; an atom is an instance of the chemical element). The atomic number of a chemical element is sometimes placed in a subscript to the left of the element symbol; for example, 6C, although this is redundant because the symbol C implies Z. Chemical elements and their reactions are the basis of chemistry.
The number of neutrons defines the element’s isotope which is represented as the element name followed by its atomic mass number (e.g. Carbon-14), so the number of neutrons may be deduced from the isotope name by subtracting the number of protons. An isotope may be abbreviated with the atomic mass number in a superscript to the left of the element symbol; for example, 14C for Carbon-14.
Atomic mass (mA) is measured in unified atomic mass units (u or amu), also known as daltons (Da) and 1u is approximately the mass of a nucleon. For various reasons, 1u is defined more strictly as
1/12th of the mass of a Carbon-12 atom. Relative atomic mass (Ar) is the weighted average mass of a set of atoms. Standard atomic weight (Ar,std) is Ar on Earth, reflecting the weighted average of isotope masses of an element on Earth.
If an atom has an equal number of protons and electrons, then it is electrically neutral. If an atom has an unequal number of protons and electrons, then it is ionized. If an atom has more protons than electrons, then it’s positively electrically charged, called a cation, and symbolized as E+ (where E is the element symbol, discussed below). If an atom has more electrons than protons, then it’s negatively electrically charged, called an anion, and symbolized as E-. If the ion is more than one electron away from the neutral element, then the + or - is preceded by that number, e.g. E2+.
The Periodic Table of Elements
There are 118 known chemical elements, 92 of which have been observed naturally, the rest synthesized, and only four elements (Carbon, Oxygen, Hydrogen, and Nitrogen) make up about 96% of a human’s mass, with another 21 required in small amounts (mostly Calcium, Phosphorus, Potassium, Sulfur, Sodium, Chlorine, and Magnesium) (Campbell & Reece, 2002, pg. 27).
The periodic table is a way to organize and understand the chemical elements based on observed patterns. The elements are ordered by atomic number Z from left to right, and starting again at the left when going down.
Within each square, the atomic number is at the top, followed by the element’s symbol, followed by the element name, followed by its standard atomic weight (Ar,std):
The main reason for describing elements in such a way has to do with electron configuration patterns and the behaviors they may cause (described later).
Rows are called periods and describe a new electron shell which accumulates on top of any previous periods’ shells. This outermost shell is known as the valence shell which generally contains the valence electrons that may be reactive (some ionized elements complicate this picture since they contain one less electron and effectively drop down a shell - for example, the Lithium cation - but the table is a conceptual starting point from neutral atoms).
Columns are called groups and generally group by the number of valence electrons, and thus generally group by similar behavior.
Each electron shell in an atom has a maximum number of electrons (
2×ShellNumber2) before the next shell starts. Each shell has a distinct energy level and is broken down into subshells which have a maximum number of electrons before the next subshell starts. Each subshell is broken down into orbitals of up to 2 electrons each. Additional electrons fill empty orbitals in a sub-shell before pairing up with an electron in an existing orbital. A full orbital is called an electron pair. A lone pair of electrons is a filled valence orbital of two electrons that are not shared with another atom (sometimes noted with a
:). Electron configurations are represented by the accumulation of subshells up to the total number of electrons, with each subshell described by the shell number, followed by the subshell name, followed by the number of electrons in that subshell in a superscript. The subshell names are:
- s: For groups 1 and 2 (or group 18 for Helium) only, at most 2 electrons.
- p: Starting at period 2, for groups 13-18 only, at most 6 electrons.
- d: Starting at period 4, for groups 3-12 only, at most 10 electrons.
- f: Starting at period 6, in between groups 3 and 4 only, at most 14 electrons.
Examples of electron configurations for the first 11 neutral elements:
Hydrogen (1): 1s1 Helium (2): 1s2 Lithium (3): 1s2 2s1 Beryllium (4): 1s2 2s2 Boron (5): 1s2 2s2 2p1 Carbon (6): 1s2 2s2 2p2 Nitrogen (7): 1s2 2s2 2p3 Oxygen (8): 1s2 2s2 2p4 Florine (9): 1s2 2s2 2p5 Neon (10): 1s2 2s2 2p6 Sodium (11): 1s2 2s2 2p6 3s1 [...]
Instead of writing the full details of long electron configurations, a common practice is to start with the previous group 18 element in [brackets] followed by the rest of the element’s electron configuration. For example, Sodium (element #11) may be written as [Ne] 3s1.
When d and f orbitals are filled, they backfill the previous shell. For example, Scandium’s (element #21) electron configuration is [Ar] 4s2 3d1.
The electrons that tend to cause an atom to chemically react are those electrons with the highest energies (the farthest distances from the nucleus) and are usually those in the valence shell. The exceptions are elements with d or f orbitals because even though those backfill the previous, non-valence shell, they may have higher energies than the s orbital electrons of the valence shell; however, these energies decrease moving right on a period, so the number of these valence electrons is limited.
Generally, atoms tend to be chemically reactive when their valence electrons do not complete their valence shell and thus the atom is unstable (high energy). For example, Hydrogen is unstable because it wants one more electron to complete its valence shell. All elements above period 1 generally want 8 electrons in their valance shell and this heuristic is called the octet rule. Atoms tend to gain, shed, or share pairs of electrons as needed to reach a full and stable (low energy) set of valence electrons. This is one of the most important aspects of chemistry and means that groups have generally similar behaviors since they’re generally grouped by the number of valence electrons. There are 18 numbered and 6 named groups for convenience:
- Group #1: Alkali metals - Highly reactive because they want to lose an electron to drop to the previous period’s full set of valence electrons.
- Group #2: Alkaline earth metals - Somewhat reactive because they want to lose two electrons to drop to the previous period’s full set of valence electrons.
- Group #15: Pnictogens
- Group #16: Chalcogens - Somewhat reactive because they want to gain two electrons or share pairs of electrons to fill their set of valance electrons.
- Group #17: Halogens - Highly reactive because they want to gain an electron or share a pair of electrons to fill their set of of valence electrons.
- Group #18: Noble gases - Generally not chemically reactive (inert) because the set of valence electrons is full.
The radius of an atom increases from top to bottom as electron shells are added; however, moving left to right, the radius of an atom decreases as the additional protons draw in the additional electrons.
Ionization energy (or cationization energy) is the amount of energy needed to remove an electron from an element (and form a cation). Energy is required to remove the electron because the electron is attracted to its element’s proton(s). Following from the octet rule, ionization energy is lowest on the left of the table because those elements want to lose electron(s) to achieve a full set of valence electrons, and ionization energy generally increases from left to right as additional protons add more pull to the electrons. Ionization energy decreases from top to bottom because valence electrons are farther from the protons (to which they’re attracted) and thus the electrons are easier to peel off.
Electron affinity (or anionization energy) is either:
- The amount of energy released when an electron is added to an incomplete set of valence electrons. Energy is released because any time an electron drops into a new orbital, it causes the release of electromagnetic radiation energy in the form of a photon. Or,
- The amount of energy spent adding an electron to create a new subshell or adding an electron to a subshell which only has a single electron in each of its orbital pairs (e.g. Nitrogen). Energy is needed to overcome the last subshell’s stability.
Generally, electron affinity increases from left to right (except for those with stable last subshells) because right-most elements want additional electrons to achieve stability.
Electronegativity is the tendency of an atom to attract electrons to its valence shell (closely related to electron affinity). It follows from the octet rule that atoms increase in electronegativity from left to right. Electronegativity decreases from top to bottom because the valence shell is farther away from the positively charged nucleus and thus there’s less pull to bring in additional electrons. Electropositivity is the opposite of electronegativity.
Most elements are metals (although there are only two named metal groups): they are toward the left side of the periodic table (with the exception of Hydrogen), have low ionization energies, low electron affinity, are highly electrically conductive, ductile, and generally solid at standard temperature.
Nonmetals are the opposite of metals: they are toward the right side of the table, have high ionization energies, high electron affinities, are not very electrically conductive, and they are often gases (e.g. Hydrogen, Helium, etc.), although some are brittle solids (e.g. Carbon).
There are a handful of Metalloids which have properties of both metals and nonmetals and run down a diagonal in the p-block (e.g. Boron, Silicon, etc.).
A transition metal is any element with a partially filled d sub-shell (groups 3-11).
In summary, although with various exceptions, the broad trends of atomic size, ionization energy, electron affinity and metallic character may be visualized as:
Big History (Continued)
After the Big Bang, the universe was mostly made of hydrogen, helium, and lithium (the first three elements), some of which combined into plasma stars held together by gravity. At high enough temperatures inside stars, hydrogen atoms undergo nuclear fusion where nuclei (and their protons and neutrons) combine to produce helium. As the hydrogen is used up to create helium, the star’s temperature rises and allows for the fusion of helium into carbon and oxygen and some neon and heavier elements. As the star temperature continues to increase, nuclear fusion produces elements up to iron. The remaining elements in the universe are produced by certain stars’ s-processes and stars’ explosions in a supernova or collapses into a black hole or neutron star.
Two or more atoms may be held together by chemical bond(s). The major types of bonds are:
- A covalent bond defines a molecule and occurs when atoms share one or more pairs of electrons in their valence shells. For example, Hydrogen wants to gain an electron and Oxygen wants to gain two electrons, so it’s common for two Hydrogens to each share their electron with one of the valence electrons in the Oxygen to form an Oxygen bonded with two Hydrogens, or water. A molecule is considered a chemical compound if it’s made of more than one type of element. The strongest form of a covalent bond is a sigma bond (σ bond). Another, weaker form of a covalent bond is a pi bond (π bond).
- A coordinate covalent bond (or dative bond) defines a covalent bond in which one atom donates both electrons to the bond. Such bonds are part of a coordination complex. An example is a metal cation which has a coordinate covalent bond to an atom with a lone pair (also called a ligand).
- An ionic bond defines an ionic compound or salt when one atom transfers electron(s) to another, creating a cation and anion, which may then cause electrostatic attraction of the oppositely charged ions. For example, the alkali metal Sodium (Na) wants to lose an electron, and the halogen Chlorine (Cl) wants to gain an electron, so Na may give its electron to Cl, thus making Na+ and Cl- and then those two ions may bond due to the electrostatic force, forming NaCl (otherwise known as table salt).
- A metallic bond amongst positively charged metal cations in a sea of shared electrons.
A chemical substance is a set of one or more elements, molecules or compounds of the same composition (i.e. “pure”). A substance cannot be separated through physical means other than breaking chemical bonds (interpret the venn diagram by the location of the bond arrows):
A mixture is a combination of different substances which cannot be separated through physical means. Therefore, an object is either a substance or a mixture. A mixture is homogenous if its substances have the same proportions throughout (e.g. air), or otherwise heterogeneous.
A molecular entity is a single instance of a part or whole of a molecule, such as an atom, ion, or molecule. A set of identical molecular entities (in other words, a class of molecular entities) is a chemical species. For example, a single water molecule is a molecular entity, but all the instances of a water molecule in some context is a chemical species. Relatedly, an elementary entity is similar to a molecular entity but may also be an electron, particle, or group of particles.
A free radical is a substance with at least one unpaired valence electron (e.g. a halogen), making it highly chemically reactive.
A substance may be described in many ways, all of which may represent either an entity or a species:
- A molecular formula describes the number of atoms of each element in each molecule in a subscript to the right of the element symbol (or 1 if the subscript is omitted). For example, the molecule H2O represents two Hydrogen atoms bonded with one Oxygen atom.
- An empirical formula is the molecular formula with the ratios of elements reduced to the simplest form. For example, a molecule of Benzene has a molecular formula of C6H6, but the empirical formula is CH.
- A trivial or retained name; for example, Water represents H2O (in the case of water, the molecular and empirical formulas are the same).
- Various structural formulas that describe the two-dimensional (2D) or three-dimensional (3D) structure of the molecule. For example, a covalent bond is represented with a long dash (–) such as
H–O–Hfor H2O. As another example, a double covalent bond is represented with a double dash (=) such as
- Wedge-hash diagrams (or wedge-dash diagrams or Natta projections): The filled wedge is projected toward the viewer (in front of the page). The dashed wedge is projected away from the viewer (behind the page).
- Fischer projections
- Unspecified stereochemistry: Wavy single bonds are unknown or unspecified stereochemistry. Stereochemistry is the study of substances with the same formula but with different positions of atoms in space that cannot be rotated around a single bond to match each other (discussed in detail later).
- Skeletal formulas
- An ionic compound name (or systematic name) with a set of element names with any cation first; for example, Sodium Chloride. If the cation is a transition metal, then it may be followed by roman numerals in parentheses (type-I compounds, type-II compounds or type-III compounds) which represents the net positive charge of that cation; for example, Iron(III) Oxide is Fe2O3 because the III means that Iron is Fe3+ and gave 3 extra electrons, and since each Oxygen atom needs two electrons, there should be two Iron atoms, making 6 extra electrons, which means there are three Oxygen atoms, each taking 2 of those 6 extra electrons.
When substances of different electronegativities are bonded, electrons will tend to be drawn closer to the atoms with higher electronegativities, thus leading to positive or negative partial charges (δ+ or δ-, respectively), creating a polarized (or polar) bond. For example, in the water molecule, Oxygen is more electronegative than the Hydrogens, so there’s a partial negative charge at the end of the Oxygen away from the two Hydrogens, and partial positive charges on the opposite ends of the Hydrogens:
Oxidation states are a conceptualizatized simplification of covalent bonds as ionic bonds. Oxidation states assume full hogging of electrons in a polar molecule. For example, in water, Oxygen would have an oxidation state of -2 since it hogs both Hydrogens’ electrons, and each Hydrogen would have an oxidation state of +1.
Reduction is the process of gaining electrons. Oxidation is the process of losing electrons. The atom losing the electrons is said to be oxidized (even if by something other than Oxygen!), i.e. losing electrons, and the atoms doing the oxidation are said to be reduced by the other elements, i.e. reducing the electrons of the oxidized atoms. This is called a redox (reduction and oxidation) reaction because both occur.
Intermolecular forces describe attraction and repulsion forces between substances and are relatively weaker than intramolecular forces:
- Dipole-dipole forces occur when the opposite partial charges of the dipoles of two polar molecules (or different parts of a large molecule) attract each other. For example, two water molecules have a dipole-dipole interaction between the partial negative end of the Oxygen of one molecule and the partial positives of the Hydrogens of the other molecule. When this interaction occurs with Hydrogren, this may be called Hydrogen bonding since it’s the strongest form of dipole-dipole interaction.
- Van der Waals forces
- London dispersion forces occur transiently at short distances as electrons happen to be in parts of their orbitals which create partial charges and create temporary dipole-dipole interactions.
When an element is anionized with one additional electron, the suffix of its name changes to -ide. The suffix -ate means more, and -ite means fewer. The prefix per- means one more, and hypo- means one fewer.
In general, matter exists in one of four states:
- Solid: Matter which has fixed volume and fixed shape, with its components close together and fixed in place. In equations, solid substances may be suffixed with (s).
- Liquid: Matter which has fixed volume and variable shape to fit its container, with its components close together but not fixed in place. In equations, liquid substances may be suffixed with (l).
- Gas: Matter which has variable volume and variable shape, both to fit its container, and its components are neither close together nor fixed in place. In equations, gas substances may be suffixed with (g).
- Plasma: Matter which has variable volume and variable shape, but also contains a large number of ions or electrons moving freely.
A mole (mol) is defined as the number of elementary entities in a substance as there are atoms in 12g of 12C. There are ~6.022140857×1023 (Avogadro constant) atoms in 12g of 12C. Therefore,
1 mol 12C = 12g.
Since 1u is 1/12th the mass of one 12C atom, then
1 mol 1u = 1g; therefore,
1 mol of isotope mAE ~= mA grams. If a substance doesn’t refer to a particular isotope, then standard atomic weight (Ar,std) is generally used instead of mA. Molecular mass is simply the sum of atomic masses of a molecule.
For example, Ar,std of H is 1.008, and Ar,std of O is 15.999, so
1 mol H2O ~= (1.008×2)g + 15.999g ~= 18.015g. The molar mass (M) of a substance is simply this relationship in terms of g/mol, so M(H2O) ~= 18.015 g/mol.
A chemical reaction occurs any time a chemical bond is created or broken. A chemical equation describes a reaction with reagents (or reactants) on the left-hand side of the equation which yields (→) the product(s) of the chemical reaction on the right-hand side of the equation (an adduct is a type of product where there are no products other than the adduct product). For example, the chemical equation for the reaction of molecular hydrogen and molecular oxygen plus input energy (E) yields water and output energy E’:
2H2 + O2 + E → 2H2O + E'
An endothermic reaction occurs when the system absorbs heat energy from its surroundings (often represented as an additional heat, light, electricity, sound, etc. reagent). An exothermic reaction occurs when the system releases energy (often represented as an additional heat, light, electricity, sound, etc. product).
Reactions that don’t require input energy are called spontaneous or exergonic.
Reactions that require input energy are endergonic:
Endergonic ⎪ E⎪ n⎪ ____ e⎪ / \________ r⎪ / products g⎪________/ y⎪reagents ⎪________________________ Time
A reagent which is attracted to electrons is called an electrophile (i.e. a lewis acid). A nucleophile is a reagent which donates electrons (i.e. a lewis base) to an electrophile.
A chemical equation must be balanced using stoichiometry because the total number of atoms of each element must be the same on both sides. A stoichiometric coefficient to the left of a substance represents the number of substances or moles of that substance in the chemical reaction.
Chemical reactions may also be reversible and go in both directions, signified by a double arrow (⇌), always tending towards reaching dynamic chemical equilibrium where the rates of reactions in both directions are equal. For example:
HCO3- + H+ ⇌ H2CO3
Given a balanced chemical equation and the relationship of moles to atomic weights, and a mass of one substance, composition stoichiometry may be used to determine the mass of other substances in the equation. For example, in the following equation:
Fe2O3 + 2Al → Al2O3 + 2Fe
If we know there are 85g of Fe2O3, since each of those molecules is about 160u, Fe2O3 is about 160g/mol, so first convert 85g to moles using dimensional analysis by multiplying by 1mol/160g to cancel the g, and therefore, 85g divided by 160g/mol is 0.53 mols. Since there are two moles of Al for every 1 mole of Fe2O3,
2×0.53 = 1.06 mols of Al would be needed to react with all of the Fe2O3. Given Al’s atomic weight is about 27u, 1 mole is about 27g, so
1.06×27 = 28.62g of Al would be needed to react with all of the Fe2O3.
Le Chatelier’s principle asserts that the reaction tends towards equilibrium even after any changes to concentration, temperature, volume, or pressure (thus why it’s called dynamic). If the temperature increases in an endothermic reaction, the system will favor the forward reaction. If the temperature decreases in an exothermic reation, the forward reaction will be favored to match the previous temperature. If the concentration of a substance increases, the reaction will favor that side of the reaction; otherwise, it will favor the other side. If volume decreases, concentrations will increase, so the reaction will favor the side that generates more total moles (i.e. creating more of the side that’s more likely to react). Conversely, if volume increases, concentrations will decrease, so the reaction will favor the side that has fewer total moles (i.e. creating more of the side that’s less likely to react).
Liquid water is a particularly imporant molecule. Since it’s a polar molecule, it has a tendency to create hydrogen bonds with nearby water molecules. This phenomenon is called cohesion. Water also has a greater surface tension (how difficult it is to stretch or break the surface of a liquid).
A calorie (cal) is the amount of heat energy it takes to raise the temperature of 1g of water by 1°C (conversely, the amount of heat energy released when 1g of water is cooled by 1°C). 1 calorie also equals 4.184 J (conversely,
1 J = 0.239 cal). Therefore, the specific heat of water is 1 cal per 1g per °C.
Relative to other substances, water has a high specific heat, meaning that it will change its temperature less when it absorbs or loses an amount of heat.
Water’s polarity makes it a common solvent.
Substances that are ionic or polar are hydrophilic meaning that they have a high affinity for water and thus increase their solubility in water. Substances that are non-ionic and non-polar are hydrophobic meaning that they have a low affinity for water and thus decrease their solubility in water. Substances that are amphipathic have both hydrophobic and hydrophilic regions.
Acids and Bases
Protium is the most common isotope of Hydrogen on Earth with one proton and zero neutrons (1H). Deuterium is an isotope of Hydrogen with one proton and one neutron (2H). A hydron is a cationic Hydrogen with 0 electrons (1 or 2H+). A hydride is an anionic Hydrogen. A protium hydron (1H+) is often just called a proton since it’s just a proton with no neutrons nor electrons.
A solution is a homogenous mixture with solute(s) dissolved into a solvent (the largest proportion substance). A solvent is often a liquid and if such a liquid is water, the solution is referred to as an aqueous solution (in equations, sometimes denoted with
(aq)). A suspension is a heterogenous mixture where one substance eventually settles, whereas colloids are heterogenous mixtures that don’t settle (emulsions which are colloids of liquids).
A litre (L) is a unit of volume equal to
1m3/1000. 1 litre of water is approximately 1 kg under standard conditions (e.g. 25°C).
The amount of a substance is usually measured in mass, moles or volume (e.g. 1g of salt in water). The concentration of substance A mixed into substance B is the amount of substance A divided by the total volume of substance B. Molarity (or molar concentration) (M) is the number of moles of a solute per liter of solution, abbreviated with square brackets around the solute (e.g. [Cl-]). Molality is the number of moles of solute per kg of solvent.
In a reversible chemical reaction such as
aA + bB ⇌ cC + dD, the equilibrium constant (keq [or kc for concentration]) equals ([C]c × [D]d) / ([A]a × [B]b), ignoring the solvent and any solids, where the a, b, c, and d coefficients are the mole ratios. Given that molarity depends on temperature, keq is a function of temperature. This constant describes the relative proportions of concentrations of reagents and products at equilibrium.
Hydroxide is the anionic molecule OH-. Hydronium is the cationic molecule H3O+.
There are multiple ways to define acids and bases:
- A Brønsted-Lowry acid is a substance capable of donating a proton. A Brønsted-Lowry base is a substance with a lone pair of electrons capable of accepting a proton (or indirectly, a substance that dissociates into hydroxide ions which then accepts the protons). A conjugate acid is the product that accepted protons (because it can then give that proton back in the reverse direction), and a conjugate base is the product that donated protons. For example, water is amphoteric meaning it can be either an acid or a base.
- A Lewis acid is a substance that has an empty electron orbital and may accept an electron pair. For example,
H+has an empty electron orbital. A Lewis base is a substance that has a lone pair of electrons not involved in bonding that it may donate.
- An Arrhenius acid is a substance that tends to react in an aqueous solution to increase the concentration of protons, or, more commonly, hydronium. For example,
HCl(aq) + H2O(l) ⇌ Cl-(aq) + H3O+(aq). An Arrhenius base is a substance that tends to react in an aqueous solution to increase the concentration of hydroxide. For example,
NaOH(aq) + H2O(l) → Na+(aq) + H3O+(aq).
The Arrhenius definition is limited to aqueous solutions whereas the Brønsted-Lowry and Lewis definitions are more general.
A substance that is basic is also sometimes referred to as alkaline or an alkali.
A strong acid or base dissociates completely in aqueous solution (→), whereas a weak acid or base dissociates partially (⇌).
A substance that is acidic in a neutral solution will have an overall negative charge. A substance that is basic in a neutral solution will have an overall positive charge.
A small proportion of water molecules in liquid water will tend to autoionize and form hydronium and hydroxide ions (
H2O + H2O ⇌ H3O+ + OH−). For pure water, at the equilibrium point under standard conditions (e.g. 25°C), [H+] = [H3O-] = 10-7 M.
The pH scale is a way to describe how acidic or basic a mixture is, and it’s just the cologarithm of molarity:
pX = -log [X]. pH is the cologarithm of H+ and describes how acidic the mixture is. Conversely, pOH is the cologarithm of OH- and describes how basic the mixture is. When
pH = pOH (in other words,
[H+] = [OH-]), the mixture is described as neutral. Mixtures with
[H+] > [OH-] are acidic. Mixtures with
[H+] < [OH-] are basic. In any mixture,
[H+]×[OH-] = 10-14 (in other words,
pH + pOH = 14). As the molarity of protons increases (e.g. stomach acid is between ~10-1.5 to ~10-3.5), the pH value decreases and thus acidity increases. Therefore, mixtures with a pH below 7 are acidic and basic above 7.
A buffer is a substance that minimizes changes to pH by accepting protons when in excess and donating protons when in deficit. Generally, a buffer is a molecule with both an acid and a base, and its equilibrium constant controls its behavior in buffering the mixture. For example, carbonic acid yields a bicarbonate ion and a proton:
H2CO3 ⇌ HCO3- + H+.
The size and structure of a substance is key to the way it functions because of the implied likelihoods of the structure on chemical reactions with other substances. A commonly used phrase is that structure implies function.
Although substance structures are often described by molecular, empirical, or two-dimensional formulas for simplicity, actual substance structures are three-dimensional.
Isomers are substances that have the some molecular formula but different structures.
Structural isomers differ in their covalent arrangements. For example, butane (C4H8) is:
H H H H | | | | H–C–C–C–C–H | | | | H H H H
Whereas isobutane (also C4H8) is:
H–C–H H ⎪ H | ⎪ | H–C–C–C–H | | | H H H
Geometric isomers (or cis-trans isomers) have the same covalent arrangements but differ in spatial arrangements. For example, the following two molecules are geometric isomers of each other. If the Xs are on the same side, it’s a cis isomer, and if they’re on opposite sides, a trans isomer:
H H \ / C=C / \ X X H X \ / C=C / \ X H
A chiral substance may have one of two, mirror-image configurations, each of which is called an enantiomer which are geometric isomers which cannot be superimposed on each other and this can have functional implications. The two forms are designated L and D isomers from the Latin for left and right (levo and dextro). R and S configurations can describe substances with multiple chiral carbons.
Ammonia is NH3. It has three covalent bonds with Hydrogen which leaves one lone electron pair (because there are 5 total valence electrons: 2s2 + 2s3). This lone electron pair makes it a base because it can accept protons. Ammonia is also polar because Nitrogen is much more electronegative than Hydrogen, but also because the shape of Ammonia is pyramidal. If it was planar, the dipoles would cancel out, but the lone pair pushes the other Hydrogens out into a pyramidal structure.
A tetrahydride is a group 14 element (e.g. C, Si, Ge, Sn, Pb, etc.) which is saturated with hydrogens (e.g. CH4). Tetrahydrides are considered non-polar.
An organic compound is an ambiguous term that historically described a compound that contains Carbon and is related to organisms. An organism is an entity with the properties of life. Even the traditional definition of organic was loose because Carbon Dioxide (
CO2) is considered inorganic even though it contains Carbon and is produced by some organisms. In addition, some (or perhaps all) organic compounds may be produced outside of organisms. The modern definition of organic retains these ambiguities, but in general, most Carbon-based compounds are considered organic. Carbon has four unpaired valence electrons and can create bonds in four different, equally spaced directions, allowing for a wide variety of compounds.
Hydrocarbons are organic molecules composed of Carbon and Hydrogen atoms. The electronegativities between Hydrogen and Carbon are close enough (2.2 and 2.55, respectively, on the Pauling scale) that hydrocarbons are generally considered non-polar and therefore generally hydrophobic.
Hydrocarbons are either aromatic (also known as arenes) or aliphatic (also known as non-aromatic). Aromatics have cyclic (in alternating single and double covalent bonds), planar arrangements (such as Benzene) and thus are very stable and unreactive, whereas most aliphatics are acyclic and less stable. Only some aromatic substances have pleasant smells.
A saturated hydrocarbon has no double or triple bonds (i.e. it’s saturated with the maximum number of Hydrogens). An unsaturated hydrocarbon has double or triple bonds between Carbons.
It is common to omit Carbons and Hydrogens in structural formulas for simplicity. For example, the structural formula for Cyclohexane (C6H12) is a ring:
H H \ / H C H \ / \ / C C /⎪ ⎪\ H ⎪ ⎪ H H ⎪ ⎪ H \⎪ ⎪/ C C / \ / \ H C H / \ H H
This can be simplified into a simple hexagon where each vertex is an implied Carbon atom with two Hydrogens:
Similarly for structures with double covalent bonds, the ring can be simplified with double lines. For example, the structural formula for Benzene (C6H6) is a ring:
H | C / ⑊ H–C C–H ║ ⎪ H–C C–H \ ⫽ C | H
This can be simplified into the following hexagon with three double bonds where each vertex is an implied Carbon atom with one Hydrogen:
Naming Organic Compounds
Greek letters may be used instead of numbers (in the same way that some numbered lists use
a, b, c, ...). The first 10 Greek letters are (lower case followed by upper case):
- Alpha: α Α
- Beta: β Β
- Gamma: γ Γ
- Delta: δ Δ
- Epsilon: ε Ε
- Zeta: ζ Ζ
- Eta: η Η
- Theta: θ Θ
- Iota: ι Ι
The longest linear covalently bonded chain of Hydrocarbons may be called a backbone chain (or main chain or carbon skeleton) from which the rest of the molecule builds off of. Any other smaller branches (in place of Hydrogens for the branch points) are called side chains, substituents, or pendant chains. Both backbone chains and side chains may be represented with the letter R (or R’, R’’, etc. if there are multiple), meaning root or remainder.
A moiety is a named part of a molecule which is a common pattern of atoms that may be found in many other molecules.
The -yl suffix is used for free radicals or moieties (with an unpaired valence electron) that can be side chains substituting for a single Hydrogen on a backbone. The -ylidene suffix is the same but when substituting two Hydrogens on the backbone for a double bond. The -ylidyne is when substituting three Hydrogens on the backbone for a triple bond.
An alkane is an acyclic saturated hydrocarbon. An example is Methane which is CH4 and is the main component of natural gas. A straight-chain (non-branched) alkane has the suffix -ane.
An alkene (or olefin) is an unsaturated alkane with at least one C=C double bond. An alkyne is an unsaturated alkane with at least one C≡C triple bond. The position of the double or triple bond is written at the start of the name or before the -ene or -yne suffix. If there are more than one such bonds, the positions are comma separated (e.g. 2,4-pentadiene). A substance that’s both an alkene and alkyne has the suffix -enyne.
An alkyl is an alkane without one Hydrogen.
The number of Carbons in a backbone chain may be thought of as the size of the alkane and gives it a prefix:
- 1: Meth-
- 2: Eth-
- 3: Prop-
- 4: But-
- 5: Pent-
- 6: Hex-
- 7: Hept-
- 8: Oct-
- 9: Non-
- 10: Dec-
- 11: Undec-
- 12: Dodec-
- 13: Tridec-
- 14: Tetradec-
- 15: Pentadec-
- 16: Hexadec-
- 17: Heptadec-
- 18: Octadec-
- 20: Eicos-
- 22: Docos-
- 23: Tricosa-
- 24: Tetracos-
- 26: Hexacos-
- 28: Octacos-
- 30: Triacont-
- 32: Dotriacont-
- 33: Tritriacont-
- 34: Tetratriacont-
- 35: Pentatriacont-
- 40: Tetracont-
- 50: Pentacont-
- 60: Hexacont-
- 70: Heptacont-
- 80: Octacont-
- 90: Nonacont-
- 100: Hect-
An aryl is an aromatic hydrocarbon minus one Hydrogen such as a phenyl (a Benzene minus one Hydrogen).
Naming an organic compound:
- An alkane with branched groups is prefixed with the position of the Carbon where the branch occurs, followed by a prefix for the branched group, followed by the name of the alkane chain.
- If there are multiple branches, the number prefix is a comma-separated list and the branch prefix is prefixed with di-, tri-, tetra-, etc. for the number of groups.
- If the multiple branches are different types of groups, they’re ordered in alphabetical order.
- For alkenes and alkynes, the position(s) of the double or triple bonds are a comma separated list infixed in the name, and a prefix of di-, tri- etcs before the final prefix.
- Prefix cis- or trans- if needing to describe the isomerism.
The moieties of an organic molecule that are most commonly involved in chemical reactions are called functional groups. When discussing functional groups, a convention is sometimes used which describes their molecular formulas without internal bond symbols, and the group is prefixed or suffixed by a dash depending on where the functional group bonds to the rest of the compound. Common functional groups:
R–OH): Oxygen covalently bonded to a Hydrogen (
O–H) and the Oxygen is covalently bonded to the compound. The highly electronegative Oxygen causes that part of the compound to be hydrophilic. Compounds containing a hydroxyl group are called Alcohols and names are usually suffixed with -ol. This functional group is hydrophilic.
For example, Ethanol is the drug in alcoholic beverages:
H H | | H–C–C–O–H | | H H
- Carbonyl (
R–CO): Oxygen double-covalently bonded to a Carbon (
C=O) which is covalently bonded to any two other atoms.
- Acyl (
R–CO): A type of carbonyl in which R is an alkyl or aryl. If the carbonyl/acyl is at the end of a chain, the group is called an aldehyde; otherwise, it’s called a ketone. This functional group is hydrophilic.
- Methyl (
R–CH3): A Carbon single bonded to three Hydrogens. Adding such a functional group makes the molecule methylated.
- Acetyl (
R–CH3CO): An acyl covalently bonded to a methyl and the remainder.
- Carboxyl (
R–COOH): A Carbon acting (i.e. not part of the backbone) in both an Acyl and Hydroxyl. Also called Carboxylic Acids (or organic acids). The reason carboxyls tend to be acidic whereas hydroxyls tend not to be is because there are two highly electronegative Oxygens near the Hydrogen in a carboxyl. This functional group is hydrophilic.
- Amino (
R–NH2): Two Hydrogens covalently bonded to a Nitrogen (
H–N–H) which is covalently bonded to the compound. Also called Amines. An amino group often acts as a base because the Nitrogen has one unpaired valence electron willing to be shared with a Hydrogen proton. Compounds that have both amino and carboxyl functional groups are called amino acids. This functional group is hydrophilic.
- Sulfhydryl (
R–SH): A sulfhydride (Sulfer bonded to a Hydrogen) which is covalently bonded to the compound. Also called thiols. This functional group is hydrophilic.
- Phosphate (
R–OPO32-): A Phosphorus double-covalently bonded to an Oxygen and three Carbonyl side chains (
O=P–(OR)3). Also called Organophosphates (or organic phosphates). This functional group is hydrophilic and anionic. Confers the ability to react with water to release energy.
- Ester (
RCOOR′): A product of a reaction between an acid (e.g. carboxylic acid) and an alcohol in which the ester bond forms where two hydroxyl groups (one from each) are replaced with a Carbonyl that bonds the two. Usually suffixed with -oate.
The first Carbon to which a functional group is attached is called the 1-Carbon or the α-Carbon. If there are multiple functional groups, generally the 1-Carbon refers to the functional group responsible for the name of the molecule.
If functional groups are on the same side of the carbon chain, they are called cis- (often represented as E), whereas trans- are on opposed sides (often represented as Z). In general, such groups contain double bonds.
A polymer is a chain-like molecule made of repeating parts called monomers connected by covalent bonds. The process of creating polymers from monomers is called polymerization. A monomer is simply a molecule that can undergo polymerization. A homopolymer is created from just one type of monomer, whereas copolymers may be created from more than one type of monomer. Oligomers are polymers of a small, fixed number of monomers (e.g. dimers, trimers, tetramers, etc.).
Polymers (and some non-polymers) are assembled through a process called dehydration synthesis (or condensation reaction or Zimmer’s hydrogenesis) in which a Hydrogen on one end of one monomer and a Hydroxyl on the other end of the other monomer combine to form water and break away, leaving the remaining monomers to covalently bond. Dehydration synthesis requires energy input. Hydrolysis is the opposite reaction (lysis is Greak for the word break) and produces energy. The verb is to hydrolize.
A macromolecule is a large molecule, often a large polymer.
Biology is the study of life. The definition of life is controversial.
A biomolecule is a molecule that’s related to some biological process, often a macromolecule but also smaller molecules.
Carbohydrates (or saccharides) are biomacromolecules made of a hydrocarbon backbone with a Carbonyl group and some number of Hydroxyl groups. Carbohydrates may also be called sugars (sacchar is Greek for sugar) although there is also a non-biochemistry term of table sugar which refers to Sucrose, a particular type of carbohydrate, which is different from the more general biochemistry term. Carbohydrates have the suffix -ose.
Monosaccharides are carbohydrates made of a single sugar molecule with a Carbonyl and multiple Hydroxyl groups. Also called simple sugars. Sugars vary in the location of their Carbonyl group, the length of the backbone (3-7 Carbons) and their spatial arrangement. Trioses: 3 Carbon sugars. Pentoses: 5 Carbon sugars. Hexxoses: 6 Carbon sugars. If the Carbonyl is at the end of the backbone, it’s an aldose; otherwise, a ketose.
One of the most common sugars used by organisms for energy is Glucose which is a hexose aldose. It has the linear arrangement:
H O \ ⫽ C ⎪ H–C–OH ⎪ HO–C–H ⎪ H–C–OH ⎪ H–C–OH ⎪ H–C–OH ⎪ H
In aqueous solution, sugars tend to form rings. In Glucose, the Carbons are numbered starting with the top Carbon (with the Carbonyl) as 1. These numbers are used to describe where bonds occur between multiple monosaccharides. The emphasized edges indicate that we’re looking at the ring edge-on, and the other non-ring bonds lie above or below the plane of the ring. There are two ring isomers of Glucose depending on where the Hydroxyl is. The α-Glucose ring is:
CH2OH ⎪ O /‾‾‾‾‾\ H /⎪ \ H ⎪/ H \⎪ ⎪\ /⎪ HO \ OH H / OH \⎪___⎪/ ⎪ ⎪ H OH
The β-Glucose ring is:
CH2OH ⎪ O /‾‾‾‾‾\ H /⎪ \ OH ⎪/ H \⎪ ⎪\ /⎪ HO \ OH H / H \⎪___⎪/ ⎪ ⎪ H OH
Disaccharides are carbohydrates made of Tto monosaccharides covalently donded. Also called double sugars. The covalent bond is called a glycosidic linkage.
Polysaccharides are carbohydrate polymers made of multiple sugars with glycosidic linkages.
Storage polysaccharides are used by organisms to store energy for future breakdown and use. Starch is a storage polysaccharide made of α-Glucose monomers (amylose being an unbranched starch, and amylopectin being branched). Glycogen is like amylopectin but even more branched.
Structural polysaccharides are used by organisms to build strong structural material. Cellulose and Chitin are structural polysaccharides made of β-Glucose monomers (and unbranched because of the β-Glucose glycosidic bonds). Parallel strands may hydrogen-bond between the Hydrogens and Hydroxyls with other strands to form microfibrils.
Lipids are a broad category of biomolecules which are mostly hydrocarbons and are only soluble in non-polar solvents (i.e. hydrophobic). Lipids are not polmers because although they form with dehydration synthesis, any repeating subunits of lipids are not directly bonded.
Glycerides are lipids with an ester bond between Glycerol and one or more fatty acids. The ester bond is made with dehydration synthesis.
Glycerol (or glycerine or glycerin) is an alcohol with a 3 Carbon backbone and 3 Hydroxyls:
H ⎪ H–C–OH ⎪ H–C–OH ⎪ H–C–OH ⎪ H
A fatty acid is an aliphatic hydrocarbon (either saturated or unsaturated) with a Carboxyl and with a (usually) 12- to 18 Carbon backbone:
OH O OH O \ ⫽ \ ⫽ C C \ ⎪ H–C–H C / ⫽ H–C–H C [...] [...] ⎪ / \ H H H
A zigzag form suggests the actual tetrahydral (in the case of a saturated fatty acid) or double covalent bond kinks (in the case of an unsaturated fatty acid) 3D orientations of the bonds. Saturated fatty acids can pack more tightly because of the flexibility of the single bonds compared to double bonds in unsaturated fatty acids (and thus saturated fatty acids tend to be more solid at room temperature). An oil is mostly made of unsaturated fatty acids. A hydrogenated oil means that unsaturated fatty acids were converted to saturated fatty acids. Most double bonds in fatty acids are of the cis- isomeric form but trans fatty acids are of the trans- isomeric form.
Monounsaturated fatty acids have a single double bond whereas polyunsaturated fatty acids have two or more double bonds.
A fat is a triaglycerol (or triglyceride) with one Glycerol and three fatty acids.
Phospholipids are lipids made of a Diglyceride and the third Hydroxyl of the Glycerol is bonded to a Phosphate group which has a negative charge. Another polar molecule such as Choline is sometimes bonded to the Phosphate. This forms a molecule which has a hydrophobic “tail” made of the fatty acids and a hyrophilic “head” made of the Phosphate.
When added to water, phospholipids form a bilayer with two hydrophilic sides opposite each other and then hydrophilic fatty acids in between.
Steroids are lipids made of four fused hydrocarbon rings and various functional groups:
/\ /\ | |__| /\ / \/ | | | \/ ⑊/
Waxes are similar to glycerides except that the Glycerol alcohol is replaced with a longer (12 to 32 Carbon) alcohol.
Proteins are biomacromolecule polymers made of one or more polypeptides. A polypeptide is an unbranched polymer of amino acid monomers bonded through peptide bonds.
An amino acid is a 1 Carbon backbone with an amino group, a carboxyl group, and one of 20 side chains. Amino acids are all of the L enantiomer form. Amino acids may be grouped by properties of their side chain (with single- and three-letter abbreviations):
- Glycine (G/Gly)
- Alanine (A/Ala)
- Valine (V/Val)
- Leucine (L/Leu)
- Isoleucine (I/Iso)
- Methionine (M/Met)
- Phenylalanine (F/Phe)
- Tryptophan (W/Trp)
- Proline (P/Pro)
- Serine (S/Ser)
- Threonine (T/Thr)
- Cysteine (C/Cys)
- Tyrosine (Y/Tyr)
- Asparagine (N/Asn)
- Glutamine (Q/Gln)
Hydrophilic + Acidic (negatively charged):
- Aspartic acid (D/Asp)
- Glutamic acid (E/Glu)
Hydrophilic + Basic (positively charged):
- Lysine (K/Lys)
- Arginine (R/Arg)
- Histidine (H/His)
Amino acids may bond when a Carboxyl group from one and an Amino group from another perform a dehydration synthesis to form a peptide bond. One end of the polypeptide has an amino group which is called the N-terminus and the other end has a Carboxyl group which is called the C-terminus. Generally, amino acids are written and read starting at the N-terminus. The polypeptide backbone is the bonded amino acids without taking into account the side chains.
The primary structure of a protein is its linear sequence of amino acids.
The secondary structure of a protein is driven by Hydrogen bonding between the polypeptide backbone. Within the polypeptide backbone, the Oxygen in the Carboxyl has a partial negative charge and the Hydrogen in the Amino has a partial positive charge. One secondary structure is the α helix which is a coil held together by Hydrogen bonding between every fourth amino acid. Another secondary structure is the β pleated sheet in which segments (β strands) of the polypeptide are parallel and held together by Hydrogen bonding between the strands.
The tertiary structure of a protein is driven by interactions between the amino acid side chains. As a polypeptide forms in water, amino acids with hydrophobic side chains tend to cluster in the core of the protein. Disulfide bridges may form where the Sulfurs in two Cysteine amino acids covalently bond. Ionic bonds may form between the negatively charged Oxygen in the Carboxyl of one amino acid and the positively charged NH3+ of a Lysine.
If a protein is made of more than one polypeptide, the quaternary structure of a protein is driven by the interactions between the polypeptides.
A protein’s shape depends on pH, salt concentration, temperature and other factors and it may become denatured and biologically inactivated if any of these factors change significantly. Excessively high fevers may be fatal because proteins may denature.
Nucleic acids (or polynucleotides) are biomacromolecule polymers made of nucleotide monomers. A nucleotide is made of a nucleoside and one to three Phosphate groups. A nucleoside is made of a Nitrogen-containing (nitrogenous) base and a five-Carbon pentose sugar. The Nitrogenous molecules are called bases because the Nitrogen atoms tend to take up H+ from solution, thus acting as bases. There are five Nitrogenous bases:
- Group: Pyrimidines (one ring)
- Cytosine (C)
- Thymine (T)
- Uracil (U)
- Group: Purines (two fused rings)
- Adenine (A)
- Guanine (G)
The pentose sugar is either Ribose or Deoxyribose (a Ribose without an Oxygen at the 2’-Carbon). The Carbon numbers in the sugar have a ‘ as compared to the Carbons in the Nitrogenous base as a short-hand for referencing the Carbons in a nucleoside.
Nucleotides are linked with a dehydration synthesis between a Phosphate group and the sugars (at the 3’ Hydroxyl) of two nucleotides, creating a phosphodiester linkage. These repeating linkages create the sugar-phosphate backbone. One end of the backbone has a Phosphate group attached to a 5’-Carbon of a sugar and the other end has a Hydroxyl group attached to a 3’-Carbon of a sugar, and these are referred to as the 5’ and 3’ ends, respectively.
Deoxyribonucleic acid (DNA) may only have A, C, T, or G Nitrogenous bases. Ribonucleic acid (RNA) may only have A, C, U, or G Nitrogenous bases.
DNA has two polynucleotide strands that together create a double helix structure. The two strands run in antiparallel (opposite) directions: starting from one end of the molecule, one strand starts at 5’ and the other strands starts at 3’. The sugar-phosphate backbones are on the outside of the helix, and the Nitrogenous bases are paired through Hydrogen bonds on the inside.
Adenine may only pair with Thymine (or Uracil in RNA), and Cytosine may only pair with Guanine.
RNA has a single polynucleotide strand; however, base pairing may still occur within the strand to create a particular geometric structure.
A gene is a sequence of Nitrogenous bases in a DNA molecule which provides the instructions for creating a sequence of amino acids, and thus proteins. A genome is the set of all genes in an organism.
A cell is a set of atoms encapsulated in a phospholipid bilayer membrane (or plasma membrane or cytoplasmic membrane), mostly filled with water. For example, a medium-sized E. coli cell has about 35 billion atoms and a human body has about 1027 atoms. The area inside the membrane is called intracellular and the area outside is extracellular.
An organism is a set of one or more cells.
History of Life
Oceans formed on Earth about 4.4 billion years ago. The first prokaryotes existed at least 2 to 3.5 billion years ago. The first eukaryotes existed at least 2 billion years ago. The first land-based eukaryotes existed at least 1.3 billion years ago. The split of plants and animals occurred at least 1 billion years ago.
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- Every time I “context switch” into this after some time away, I re-read the entire page before re-starting. Along with the process of writing itself, this seems to maximize retention.