Summary
Highlights
Substances are made of atoms, represented by symbols in the periodic table. Compounds contain two or more different types of atoms chemically bonded, like H2O. Chemical reactions involve atoms changing bonds. Equations must be balanced as atoms are not created or destroyed. Quick tip: balance atoms in compounds first, then elements last.
Mixtures are combinations of elements and compounds not chemically bonded, like air or salt water. Separation techniques include filtration for insoluble particles, crystallization for dissolved solids by evaporating the solvent, and distillation where gas is cooled and condensed back to liquid. Fractional distillation separates liquids with different boiling points. These are physical processes, not chemical reactions.
The three main states of matter are solid (fixed vibrating particles), liquid (touching but free-moving particles), and gas (far apart, randomly moving, highest energy particles). Gases can be compressed, unlike solids and liquids. Energy is needed to overcome electrostatic forces for melting or evaporating, without breaking bonds. State symbols (s, l, g, aq) indicate the state of substances.
Atomic models evolved: J.J. Thomson (plum pudding model, positive charge with embedded negative electrons), Ernest Rutherford (discovered the small, positive nucleus through alpha particle scattering), Neils Bohr (electrons in shells/orbitals), and James Chadwick (neutral neutrons in the nucleus). Protons (+1), neutrons (0), and electrons (-1) have relative masses of 1, 1, and 0 (or very small) respectively.
The atomic number (bottom) is the number of protons, determining the element. Atoms have equal protons and electrons for a neutral charge. Ions are atoms that have gained or lost electrons. The mass number (top) is protons + neutrons. Isotopes are atoms of the same element with different numbers of neutrons. Relative atomic mass accounts for the average mass of all isotopes based on their abundance.
Early periodic tables ordered elements by atomic weight. Dmitri Mendeleev grouped elements by properties, leaving gaps for undiscovered elements, proving his method correct. Electron shells fill from inside out (2, 8, 8, 2 for first 20 elements).
Metals (left of staircase) donate electrons; non-metals (right) accept electrons. Group number indicates outer shell electrons. Group 1 (alkali metals) have one outer electron, becoming more reactive down the group as electrons are more easily donated due to increasing distance from nucleus. Group 7 (halogens) have seven outer electrons, becoming less reactive down the group as electron acceptance is harder. Group 0 (noble gases) are unreactive with full outer shells.
Metals form positive ions (cations) by losing electrons (e.g., Group 1 become 1+, Group 2 become 2+). Non-metals form negative ions (anions) by gaining electrons (e.g., Group 7 become 1-, Group 6 become 2-). Transition metals can form ions with varying charges (e.g., Fe2+, Fe3+). Ionic compounds (salts) are named metal ion then non-metal ion, e.g., sodium chloride.
Metallic bonding forms a lattice of ions with delocalized electrons, making metals good conductors. Ionic bonding occurs between metals and non-metals, involving electron transfer to form oppositely charged ions. Dot and cross diagrams show electron transfer. Ionic compounds form crystals with high melting/boiling points due to strong electrostatic forces. They conduct electricity when molten or dissolved as ions are mobile.
Covalent bonding occurs between non-metals by sharing electrons to achieve full outer shells. Simple molecular structures (individual molecules) have low boiling points due to weak intermolecular forces. Giant covalent structures (e.g., diamond) involve extensive covalent bonds, resulting in high hardness and melting points. Graphite is a giant covalent structure with layers and delocalized electrons, allowing electrical conductivity and slipperiness.
Carbon allotropes include diamond, graphite, graphene (single layer of graphite), and fullerenes (3D carbon structures like buckminsterfullerene and nanotubes). Nanoparticles have a huge surface-to-volume ratio, meaning fewer are needed to achieve specific purposes compared to larger particles.
Mass is conserved in chemical reactions. Relative formula mass (Mr) is the sum of relative atomic masses (Ar) in a compound. A mole is 6.02 x 10^23 particles (Avogadro's constant). One mole of a substance has a mass in grams equal to its Mr. Moles = mass (g) / Mr. Stoichiometry refers to mole ratios in balanced equations, allowing calculation of reactant/product masses.
A limiting reactant is the one completely used up in a reaction, determining the maximum product. Concentration can be expressed in g/dm³ or moles/dm³ (molar). Percentage yield compares actual product mass to theoretical maximum, while atom economy measures the efficiency of a reaction in converting reactants into useful products.
One mole of any gas occupies 24 dm³ at room temperature and pressure (RTP). The reactivity series ranks metals by their reactivity in displacing other elements. A more reactive metal displaces a less reactive metal from its compounds (e.g., zinc displacing copper from copper sulfate).
Metals less reactive than carbon can be extracted by carbon (smelting). Reduction is the loss of oxygen or gain of electrons (OIL RIG: Oxidation Is Loss, Reduction Is Gain of electrons). Oxidation is the gain of oxygen or loss of electrons. More reactive metals react with acids, producing a salt and hydrogen.
Alkalis (pH > 7) react with acids to produce a salt and water (neutralization). The pH scale is logarithmic, meaning a pH change of 1 represents a tenfold change in H+ concentration. Strong acids fully dissociate; weak acids partially dissociate. Titrations determine unknown concentrations of acids or alkalis using indicators and known concentrations.
Electrolysis involves using electricity to break down ionic compounds. In molten ionic compounds, cations (positive ions) move to the cathode (negative electrode) and are reduced, while anions (negative ions) move to the anode (positive electrode) and are oxidized. In aqueous solutions, reactivity determines which ions are discharged at the electrodes. Less reactive ions are discharged at the cathode; halides are discharged at the anode if present, otherwise oxygen from hydroxide ions.
Energy is needed to break bonds and released when bonds form. Exothermic reactions release net energy (temperature increase), while endothermic reactions absorb net energy (temperature decrease). Energy profile diagrams show potential energy changes and activation energy. Bond energy calculations can determine the overall energy change of a reaction.
Cells (batteries) produce a potential difference from chemical reactions. Non-rechargeable batteries stop when reactants are used up. Rechargeable batteries can be reversed by an external current. Hydrogen fuel cells produce electricity by reacting hydrogen and oxygen, forming water.