Chemistry Fundamentals: Atoms, Bonding, and Reactions
Atoms and Elements
Atoms are the fundamental building blocks of matter.
Atoms consist of a core (nucleus) and electrons.
The core contains protons and neutrons.
The number of protons determines the element.
Water is a compound made of hydrogen (H) and oxygen (O).
Quantum mechanics describes atoms differently from the simple core-electron model.
Electron Shells and Valence Electrons
Atoms have multiple electron shells.
Valence electrons are those in the outermost shell.
The behaviour of valence electrons governs most of chemistry.
Periodic Table
Elements are organized in the periodic table.
Elements in the same column (group) have the same number of valence electrons.
For main groups, the number of valence electrons equals the group number (1-8), except for helium (He), which has only 2 but behaves like a noble gas.
Transition metals do not follow a straightforward pattern regarding valence electrons.
Elements with similar valence electron numbers exhibit similar chemical behaviour.
Alkali metals (Group 1, excluding hydrogen) have one valence electron, are shiny, soft, and reactive.
Elements in the same row (period) have the same number of electron shells.
Mass increases from left to right across the periodic table with each element gaining protons, electrons, and neutrons.
Isotopes are variants of an element with differing numbers of neutrons.
Many isotopes are unstable and decay, releasing ionizing radiation, which is harmful.
Ions
An atom with an equal number of electrons and protons is neutral (no charge).
If there are more electrons than protons, the atom has a negative charge (anion).
Fewer electrons than protons result in a positive charge (cation).
Charged atoms are called ions.
Periodic Table Information
Each cell in the periodic table provides:
Element name and symbol.
Number of protons (atomic number), which equals the total number of electrons in a neutral atom.
Atomic mass (combined mass of protons and neutrons).
Categories of Elements
Metals are on the left side of the table.
Nonmetals, mostly gases, are on the right side.
Semimetals (metalloids) lie along the dividing line and have intermediate properties.
Molecules and Compounds
A molecule consists of two or more bonded atoms.
A compound contains at least two different elements.
The properties of a compound can be very different from those of its constituent elements.
Example: Sodium (explosive metal) + Chlorine (toxic gas) = Table Salt (stable compound)
Molecular Formulas and Isomers
Molecular formula indicates the number of each atom type in a molecule.
Isomers are molecules with the same molecular formula but different structures and properties.
Example: Graphite and diamond are both forms of carbon but have distinct properties due to their structure.
Lewis Dot Structures
Show valence electrons and bonds as dots and lines.
Help understand why atoms bond.
Chemical Bonding and Energy
Systems tend towards states of lower energy.
Atoms achieve lower energy with a full outer electron shell (typically eight electrons, or two for hydrogen and helium).
Noble gases have full outer shells and are thus unreactive.
Atoms share electrons to achieve a full outer shell, forming a covalent bond.
Think of it like a ball rolling downhill, decreasing it's potential energy.
Sharing of electrons in a covalent bond is caused by the positively charged nucleus of one atom tugging on electrons of another atom.
The strength of this pull is called electronegativity.
Electronegativity increases from bottom left to top right on the periodic table, with fluorine (F) having the highest electronegatvity.
Ionic Bonds
Occur when the electronegativity difference between two atoms is greater than approximately 1.7.
Example: Sodium chloride (NaCl).
Chlorine strongly attracts an electron from sodium, and sodium readily gives up an electron.
Sodium becomes a cation (positive ion), and chlorine becomes an anion (negative ion).
Salts are formed from metals and nonmetals bonding ionically, creating a grid of ions.
Metallic Bonds
Occur in pure metals.
Consist of a grid of positively charged nuclei surrounded by freely moving (delocalized) electrons.
Delocalized electrons account for metals' conductivity of electricity and heat, as well as malleability.
Covalent Bonds (Polar vs. Nonpolar)
Nonpolar covalent bonds form when electronegativity difference is less than about 0.5; electrons are shared equally.
Polar Covalent Bonds:
Electronegativity difference is between 0.5 and 1.7.
Electrons are unequally shared, creating partial charges.
Example: Water (H₂O) - Oxygen pulls electrons more strongly than hydrogen, resulting in partial negative charge on oxygen and partial positive charges on hydrogen.
Electric Dipoles and Intermolecular Forces (IMFs)
An electric dipole is the presence of two poles with opposite charges.
Permanent dipole molecules interact, arranging themselves with opposite charges aligned.
Intermolecular forces (IMFs) are forces acting between molecules.
Hydrogen Bonds:
Occur when hydrogen is bonded to a highly electronegative atom (fluorine, oxygen, or nitrogen), creating strong dipoles.
Van der Waals Forces:
Temporary dipoles formed by chance electron distribution induce dipoles in neighbouring particles.
Weaker than hydrogen bonds, but still significant.
Solvents and Polarity
Water (polar) can dissolve polar molecules by interacting with their charges.
Water cannot dissolve nonpolar molecules.
"Like dissolves like".
Surfactants (e.g., soap) have polar heads and nonpolar tails, allowing them to surround nonpolar substances (e.g., fats) in water, forming micelles that transport dirt away.
Bond and Force Strength Ranking
Ionic bonds > Covalent bonds > Metallic bonds > Hydrogen bonds > Van der Waals forces.
States of Matter
Solid:
Tightly packed particles in a fixed arrangement, only wiggling.
Liquid:
Particles move freely but remain in a fixed volume due to intermolecular forces.
Gas:
Particles have enough energy to move freely and fill any available volume.
Temperature and Entropy
Temperature is the average kinetic energy of particles in a system.
Entropy is the measure of disorder.
Solids exist at low temperature and/or high pressure (low entropy).
Gases exist at high temperature and/or low pressure (high entropy).
Strong bonds, like ionic bonds, result in high melting points.
Plasma
A fourth state of matter: ionized gas at very high temperatures or electric potential.
Emission Spectrum
Gas can be ionized in a tube with a very high voltage.
Collisions of the ions with other particles makes their electrons move to a higher energy state.
Once they fall back down, the difference in energy is released as light.
The color of the light depends on the element that's used in the tube as each element has different but fixed energy levels and the difference between those determines the energy and therefore the frequency of the released light, which is what changes the color.
All possible frequencies than an element can emit are called the emission spectrum.
Mixtures
Pure substances consist of one element or compound.
Mixtures contain at least two pure substances.
Homogeneous Mixtures:
Substances mix evenly; the mixture looks the same everywhere.
Example: Saltwater (solution).
Heterogeneous Mixtures:
Have distinct regions with different substances.
Example: Sand and water (suspension).
Colloids
Colloids (e.g., milk or more precisely an emulsion) appear homogeneous but have larger particles than solutions.
Particles are evenly distributed but not fully dissolved, between solutions and suspensions.
Chemical Reactions
Explosions are rapid chemical reactions that release energy and expand quickly.
Types of Chemical Reactions:
Synthesis
Decomposition
Single replacement
Double replacement
Reactions occur to decrease energy and achieve a more stable state.
Stoichiometry describes the ratios in which chemical reactions occur.
For example, 2H2 + O2 \rightarrow 2H_2O requires twice as much hydrogen as oxygen.
The conservation of mass states that mass cannot be created or destroyed but only converted.
Balancing Chemical Equations
Ensure the same number of each type of atom on both sides of the equation.
Balance metals first, then nonmetals, and hydrogen and oxygen last.
Moles
Atomic mass in grams yields one mole (6.022 x 10^{23} particles).
Physical vs. Chemical Changes
Physical changes alter appearance without changing the substance (e.g., hammering metal).
Chemical changes alter the substances themselves and are often accompanied by bubbles, smells, or explosions.
Activation Energy and Catalysts
All chemical reactions require activation energy.
Catalysts reduce the activation energy, making reactions easier and faster, and are not consumed in the reaction.
Enthalpy
Enthalpy (H) is the internal energy or heat content of a system.
Exothermic Reactions:
Release heat; the total enthalpy is lower at the end than at the beginning (negative \Delta H).
Endothermic Reactions:
Absorb heat; the total enthalpy is higher at the end (positive \Delta H).
Gibbs Free Energy
Gibbs free energy (G) considers both enthalpy (H) and entropy (S) at a given temperature (T): \Delta G = \Delta H - T\Delta S
Exergonic Reactions:
Spontaneous reactions that release free energy (\Delta G < 0 ).
Endergonic Reactions:
Non-spontaneous reactions that require energy input (\Delta G > 0).
Temperature and entropy play roles in reaction spontaneity.
Melting ice is endothermic but can be spontaneous above 0°C due to increased entropy.
Equilibrium occurs when \Delta G = 0; reactions proceed at the same rate in both directions.
Acids and Bases
Acids donate protons (H⁺ ions).
Bases accept protons.
Conjugate acids and bases are formed when acids/bases donate/accept protons.
Amphoteric molecules can act as both acids and bases (e.g., water).
Strong acids dissociate completely into ions, producing many hydronium ions (H₃O⁺).
Weak acids dissociate less, resulting in lower hydronium ion concentrations.
pH Scale
pH measures the strength of an acid based on the hydronium ion concentration: pH = -log[H_3O^+]
Each pH unit represents a tenfold change in concentration.
pH of 7 is neutral.
pH < 7 is acidic.
pH > 7 is basic.
pOH measures basicity.
pH + pOH = 14
Neutralization occurs when strong acids and bases combine to form water and a neutral salt.
Redox Reactions
Reduction-oxidation (redox) reactions involve the transfer of electrons.
Oxidation is the loss of electrons.
Reduction is the gain of electrons.
Oxidizing agents (oxidants) cause oxidation and are themselves reduced.
Reducing agents (reductants) cause reduction and are themselves oxidized.
Oxidation Numbers:
Imaginary charges assigned to atoms in a molecule.
Hydrogen is mostly plus one.
Oxygen is mostly negative two.
Halogens are mostly negative one.
Single elements are always zero.
The number of all atoms in a molecule always have to add up to the molecule's charge.
Redox reactions change the oxidation numbers of elements.
Redox reactions in acidic or basic solutions can be balanced using ions (H⁺ or OH⁻) and water.
Quantum Numbers and Electron Configuration
Electrons are described by four quantum numbers: n, l, mₗ, and mₛ.
n: principal quantum number (shell).
l: azimuthal quantum number (subshell shape).
mₗ: magnetic quantum number (orbital orientation).
mₛ: spin quantum number (electron spin, +1/2 or -1/2).
Orbitals are three-dimensional regions where electrons are likely to be found.
Described by Schrodinger's equation, giving a probabilistic wave function.
Subshells: s, p, d, f.
l determines the subshell; n, l, and mₗ determine the orbital.
Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
Each orbital can hold a maximum of two electrons with opposite spins.
Number of Orbitals and Electrons in Each Subshell:
s: 1 orbital, 2 electrons.
p: 3 orbitals, 6 electrons.
d: 5 orbitals, 10 electrons.
f: 7 orbitals, 14 electrons.
Shell Capacities:
Shell 1: 2 electrons.
Shell 2: 8 electrons.
Shell 3: 18 electrons.
The number of electrons a shell can hold follows the rule 2n^2
Aufbau Principle: Subshells are filled in a specific order.
Electron Configuration: Describes the arrangement of electrons in an atom's subshells and orbitals.
Condensed Electron Configuration: Uses the previous noble gas to shorten the notation.