Chemistry

Unit 1, 2

Chemistry Flashcards - Oxford T.B.

Chapter 1 (1.1, 1.2, 1.3, 1.4)

Element

Primary constituents of matter that cannot be chemically broken down into simpler substances

Compound

Atoms of different elements chemically bonded together in a fixed ratio

Mixture

Contains more than one element or compound in no fixed ratio; not chemically bonded; can be separated by physical methods

Homogeneous mixture

Evenly distributed particles (e.g. air)

Heterogeneous mixture

Unevenly distributed particles (e.g. mixture of two solids)

States of matter

Solid; liquid; gas

Solids particle arrangement

Fixed positions vibrate in lattice strong intermolecular forces

Liquids particle arrangement

Mobile close together weaker forces than solids

Gases particle arrangement

Far apart move freely negligible intermolecular forces

Sublimation

Solid to Gas directly (without melting). Example: dry ice (CO₂)

Deposition

Gas to Solid directly (without condensing). Example: snowflakes

Endothermic changes of state

Absorbs energy (solid to liquid to gas)

Exothermic changes of state

Releases energy (gas to liquid to solid)

Temperature

Measure of average kinetic energy of particles

Absolute zero

0 K (-273.15°C); particles cannot transfer kinetic energy; cannot get colder

Base SI units

kg (mass); m (length); s (time); A (electric current); K (temperature); mol (amount); cd (luminous intensity)

Mole

SI unit of amount of substance containing 6.02 × 10²³ particles (Avogadro's constant)

Kinetic molecular theory

Model to explain physical properties of SOM and changes of state

Filtration

Separates insoluble solids from liquids using filter paper

Distillation

Separates miscible liquids with different boiling points

Crystallization

Separates soluble substance from solution by evaporating solvent and cooling

Chromatography

Separates mixtures based on different affinities for mobile and stationary phases

Chemistry Flashcards - Oxford

Chapter 2 (2.1, 2.2, 2.3, 2.4)

Atom structure

Positively charged nucleus (protons + neutrons) surrounded by negatively charged electrons

Nucleons

Protons and neutrons (found in the nucleus)

Rutherford's gold foil experiment

Most alpha particles passed through (atom mostly empty); some deflected (positive nucleus); few bounced back (dense nucleus)

Proton relative mass and charge

Mass 1 charge +1

Neutron relative mass and charge

Mass 1 charge 0

Electron relative mass and charge

Mass negligible charge -1

Atomic number

Number of protons in the nucleus; equals electrons in a neutral atom

Mass number

Protons + Neutrons in the nucleus

Neutrons calculation

Mass number - Atomic number

Isotopes

Atoms of same element with different numbers of neutrons (same protons but different mass)

Relative atomic mass (Aᵣ)

Weighted average of all isotopes based on natural abundance

Aᵣ formula

Aᵣ = Σ(mass of isotope × natural abundance) / 100

Mass spectrometry

Detects relative abundance of isotopes in a sample

m/z ratio

Mass to charge ratio; species with lowest m and highest z deflected most

Cation

Positively charged ion formed when atom loses electrons

Anion

Negatively charged ion formed when atom gains electrons

Emission spectrum

Series of lines against dark background produced when excited electrons return to lower energy levels

Absorption spectrum

Series of dark lines within continuous spectrum when cold gas absorbs specific wavelengths

Quantization

Electromagnetic radiation comes in discrete packets (quanta/photons)

Bohr model

Electrons exist in discrete energy levels/stationary orbits; transitions between levels emit/absorb photons

Ground state

Lowest possible energy state of an atom (n=1 for hydrogen)

Excited state

Higher energy level (n=2; 3 ...) that is unstable; atom returns to ground state emitting photons

Maximum electrons in energy level n

2n² electrons

Atomic orbital types

s (spherical); p (dumbbell); d/ f (in order of increasing energy)

Atomic orbital

Region of space with high probability of finding an electron

Maximum electrons in s orbital

2 electrons

Maximum electrons in p sublevel

6 electrons (3 orbitals × 2 electrons each)

Maximum electrons in d sublevel

10 electrons (5 orbitals × 2 electrons each)

Pauli Exclusion Principle

Only two electrons can occupy same orbital and they must have opposite spins

Hund's Rule

Every degenerate orbital is singly occupied before any is doubly occupied; all singly occupied orbitals have same spin

Aufbau Principle

Electrons fill lowest available energy orbitals before higher energy ones

Orbital filling order

1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d < 5p < 6s < 4f < 5d < 6p

Degenerate orbitals

Orbitals with the same energy (e.g.

Chromium electron configuration exception

[Ar] 4s¹ 3d⁵ (not [Ar] 4s² 3d⁴) - half-filled d sublevel is more stable

Copper electron configuration exception

[Ar] 4s¹ 3d¹⁰ (not [Ar] 4s² 3d⁹) - full d sublevel is more stable

Ionization energy

Minimum energy required to remove an electron from a neutral atom in ground state

Ionization energy trend across a period

Increases across a period (increased nuclear charge

Ionization energy trend down a group

Decreases down a group (more shielding

Ionization energy discontinuities

1) Group 2 to 13: paired 2s electrons shield 2p electron (lower IE); 2) Group 15 to 16: half-filled p sublevel in N is more stable than paired electrons in O

Transition element

Element with a partially filled d sublevel (or forms ions with incomplete d sublevel)

Avogadro's constant (Nₐ)

6.02 × 10²³ mol⁻¹ (number of particles in one mole)

N = n × Nₐ

Number of particles = moles × Avogadro's constant

Relative molecular mass (Mᵣ)

Sum of Aᵣ values for all atoms in a molecule (no units)

Molar mass (M)

Mass of 1 mole of a substance; units = g mol⁻¹

n = m / M

Moles = mass / molar mass

Empirical formula

Simplest whole number ratio of atoms of each element in a compound

Molecular formula

Actual number of atoms of each element in a molecule

Molar concentration (c)

Amount of solute per volume of solution; units = mol dm⁻³

c = n / V

Molar concentration = moles / volume

Standard solution

Solution with a known concentration of solute

Dilution formula

c₁V₁ = c₂V₂

Avogadro's law

Equal volumes of any two gases at same temperature and pressure contain equal numbers of molecules

Gas volumes and moles

n₁/n₂ = V₁/V₂ (under same T and P)

Ideal gas assumptions

1) Constant random motion; 2) Elastic collisions; 3) Negligible molecular volume; 4) No intermolecular forces; 5) Kinetic energy ∝ Kelvin temperature

Boyle's Law

At constant T; pressure is inversely proportional to volume: pV = constant

Real gas deviation conditions

Low temperature and high pressure (intermolecular forces matter; molecular volume significant)

Ideal gas behavior conditions

Low pressure and high temperature

Molar volume of ideal gas at STP

22.7 dm³ mol⁻¹ (0°C; 100 kPa)

Combined gas law

p₁V₁/T₁ = p₂V₂/T₂

Ideal gas equation

pV = nRT (R = 8.31 J K⁻¹ mol⁻¹)

Ionic bonding

Electrostatic attraction between oppositely charged ions (cations and anions)

Electronegativity

Measure of atom's ability to attract a pair of covalently bonded electrons

Ionic bond criterion

When electronegativity difference > 1.8

Lattice enthalpy

Energy required to form gaseous ions from one mole of solid ionic lattice; endothermic process

Lattice enthalpy factors

1) Increasing ionic charge = stronger attraction = higher lattice enthalpy; 2) Increasing ionic radius = weaker attraction = lower lattice enthalpy

Ionic compound properties

High melting/boiling points; brittle; non-volatile; conduct electricity when molten or aqueous (not solid); soluble in polar solvents

Solid ionic compounds don't conduct

Ions are held in fixed lattice positions and cannot move

Ionic compounds conduct when molten/dissolved

Ions become mobile and can move to carry charge

Ionic lattice

Continuous 3D network of repeating positive and negative ions

Polyatomic ions

Ions containing several atoms covalently bonded together (e.g. NH₄⁺ NO₃⁻ CO₃²⁻)

Common polyatomic ions

NH₄⁺ (ammonium); OH⁻ (hydroxide); NO₃⁻ (nitrate); HCO₃⁻ (hydrogencarbonate); CO₃²⁻ (carbonate); SO₄²⁻ (sulfate); PO₄³⁻ (phosphate)

Ionic compound formula determination

Net charge must be zero; use criss-cross rule or bar diagram method

Covalent bond

Electrostatic attraction between a shared pair of electrons and the positively charged nuclei

Octet rule

Atoms tend to gain

Bond order

Number of shared electron pairs between two atoms (single=1; double=2; triple=3)

Bond order effects

Higher bond order = stronger and shorter bonds