chemistry

Relative atomic mass

The average mass of an atom of an element relative to 1/12 the mass of a carbon-12 atom.

Isotope

Atoms of the same element with the same number of protons but different numbers of neutrons.

Atomic number

the number of protons in the nucleus of an atom. Mass number is the sum of protons and neutrons in the nucleus.

Elements, compounds, mixtures

An element is a substance made up of only one atom.

Compounds are joined by 2 or more atoms by any type of chemical bond

Mixtures are made from 2 ions, compounds, mixed together but not chemically bonded

Mole

A different way to count the particles within a substance

Solution, solute and solvent

Components of a solution where solute is dissolved in a solvent to form a homogeneous mixture.

Empirical formula

The simplest whole number ratio of atoms of each element present in a compound.

Molecular formula

A chemical formula indicating the actual number of atoms of each element in a molecule.

Skeletal formula

A simplified organic formula showing atom arrangement without all bonds.

Displayed formula

A molecular formula showing all atoms and bonds in a molecule.

Homologous series

Organic compounds with the same functional group and similar properties.

Structural isomer

Isomers with the same molecular formula but different structures.

Stereoisomer

Isomers with the same molecular and structural formula but different has spatial arrangements.

Optical isomer

Non-superimposable mirror image stereoisomers.

Chiral centre

An atom in a molecule bonded to four different groups.

Complete and incomplete combustion

Complete combustion produces CO2 and H2O

Incomplete combustion produces CO or C.

Molecular ion peak

Peak in a mass spectrum corresponding to the molecular ion.

Fragment

A smaller molecule formed during mass spectrometry.

Isotopic abundance

Relative amount of each isotope of an element in a sample.

SN1 reaction

A two-step reaction forming a carbocation intermediate.

SN2 reaction

A one-step reaction with simultaneous nucleophile attack and leaving group departure.

SEAR reactions

Single electron transfer-aromatic substitution reactions.

Essential amino acid

Amino acid essential from the diet, not synthesized by the body.

Buffer

A solution resisting pH changes with acid or base addition.

Optimum temperature/pH

The most efficient temperature or pH for a reaction.

Dynamic equilibrium

State with constant reactant/product concentrations due to equal forward and reverse reaction rates.

Closed system

A system where neither matter nor energy can enter or leave.

Position of equilibrium

Relative reactant/product concentrations at equilibrium.

Le Chatelier's principle

System at equilibrium shifts to counteract changes in temperature, pressure, or concentration.

Catalyst

Substance increasing reaction rate without being consumed.

When a substance is used to quicken up the rate of a reaction without being used up itself

Equilibrium constant

Ratio of product to reactant concentrations at equilibrium.

Common ion effect

Suppression of weak electrolyte ionization by a strong electrolyte with a common ion.

Partial pressure

Pressure from one gas component in a mixture.

Mole fraction

Ratio of moles of a component to total moles in a mixture.

Rate of reaction

Change in concentration per unit time.

Collision theory

Reactions occur when particles collide with sufficient energy and orientation.

Activation energy

Minimum energy for a reaction to occur.

Maxwell-Boltzmann distribution

Distribution of molecule speeds in a gas at a given temperature.

Rate equation

Relates reaction rate to reactant concentrations.

Rate constant

Proportionality constant in the rate equation.

Order of reaction

Exponent of the concentration term in the rate equation.

Strong acid/base

Acid/base fully dissociating in solution.

Weak acid/base

Acid/base ions partially dissociating in solution.

pH scale

Scale indicating solution acidity or basicity between 1-14.

pKa

Negative logarithm of the acid dissociation constant (Ka).

Kw – self ionisation of water

Equilibrium constant for water autoionization.

Spectroscopy

Study of matter-electromagnetic radiation interaction.

Fingerprint region

Infrared spectrum region with unique absorption bands.

it is a region between o-1500

Carbon environment

Local chemical environment influencing NMR chemical shift.

Chemical shift

Peak displacement in NMR spectrum, measured in ppm.

Integral

Area under an NMR peak, proportional to contributing protons.

Enthalpy of combustion

this is when heat energy is given out when one mole of a substance burns completely in oxygem

Standard enthalpy

the change under standard conditions

Enthalpy

it is the total energy associated with a system at constant pressure

oxidation

loss of electrons

reduction

gain in electrons

redox

combination of reduction and oxidation

hybridisation

combination of S-orbitals and P-orbitals to form new atomic orbitals from the originals

electronegativity

power of an atom to attract electron density from inside of a covalent bond

Types of intermoleculer forces

hydrogen bonding

dipole-dipole (temporary)

london forces

Collision theory

A reaction can only take place if:

particles collide

collide with sufficient energy

collides at the correct orientation

Increasing rate of reaction

add catalyst

increase temperature

increase pressure

increase surface area

stir solution (particles with collide with each other)

ionic bonding

it is an electrostatic attraction between oppotively charged ions

hydrogen bonding

it is a weak electrostatic attraction between electrons due to an uneven distribution of electrons in some atoms

Test for alcohols

white misty fumes of HCI will be released

the fumes will become acidic and turn the universal indicator paper red

Test for haloalkanes

coloured precipitate

exited state

When an electron is given energy to move up in higher energy levels

sometimes the electron can be exited so far it can leave the atom leaving a positively charged ion

Relaxation state

when energy is released from an electron making it move in lower energy levels

Atom

the smallest particle made up on protons, electrons and neutrons

metallic bonding

A metallic bond is the electrostatic attraction between the sea of delocalsed electrons and lattice of positive ions

Equilibrium constant Kc

It is a constant which describes the ratio of products and reactants in a given equilibrium at a given temperature

Ionisation energy down a group

Atomic radius increases down a group

outer electrons are further from the nucleous

electrostatic attraction is decreased

less energy is needed to remove outer electrons

Shielding increases down a group

more repulsion between inner and outer electrons

less energy needed to remove outer electrons

Ionisation across a period

Atomic radius decreases

outer electrons are held more strongly by nucleus

needs more energy to remove outer electrons from the atom

shielding fills same outer energy level across a period

S-Orbitals

s-orbitals are like electron shells​

There is only one s-orbital in each sublevel​

They increase in size as they move up the energy levels​

They fit around each other, with the nucleus in the middle​

P-Orbitals

p-orbitals have a “dumbbell” or “figure of 8” shape​

There are three p-orbitals in each p-sublevel​

Each p-orbital points in a different direction​

They fit together to make the p-sublevel

3 pieces of evidence to disprove kekule structure of benzene

Does not react with bromine water due to a lack of discreet C=C bonds

Reduced energy of hydrogenation

all bond lenghts are the same

How to identify peaks in chemical analysis

Anything between 0-1500 is a fingerprint region which can be ignored

Polar bonds

A significant difference in electronegativity values

Lattice enthalpy

When one mole of an ionic solid is formed in its standard state from the corresponding gaseous ions

Enthalpy of atomisation

when one mole of a gaseous atom is formed from its elements under standard conditions.

Electron affinity

When one mole of a gaseous atom gains one mole of electron to form one mole of a gaseous anion

spontaneous reaction

Depends on the entropy change of its surroundings which change as energy is released or absorbed

Enthalpy

Amount of energy lost or gained by a system but does not tell you whether a reaction will occur or not

Primary in haloalkane functional groups

When a haloalkane is attached to a carbon which is attached to the end of the carbon chain

secondary in functional group

When a functional group is attached to a carbon which is attached to 2 other carbons

tertiary

When a functional group is attached to a carbon attached to 3 other carbons

Electrophile

Takes H+ ion pair

nucleophile

Donates H+ ions

2nd laws of thermodynamics

Energy is neither lost nor gained but can only be transferred to another energy source

Entropy can increase constantly inside of a closed system

how does Entropy change in different processes

Entropy increases from solid to liquid to gas

The entropy change can be predicted whether if it is positive or negative by looking at the phases of the reactants and products

Formation of a covalent bond when orbitals overlap

When 2 atoms come close together the atomic orbitals overlap. The overlap allows electrons from each atom to be shared between them. The sharing of electrons creates a bond between the atoms.