Chemistry Mock Final

Atom (in terms of subatomic particles)

The smallest particle of an element that contains electrons surrounding a nucleus containing electrons and protons

Molecule

A group of two or more atoms chemically combined to form an identifiable unit which retains the properties and composition of the substance

Subatomic Particle Masses

• Protons: 1

• Electrons: 1/1840

• Neutrons: 1

Atomic Number

number of proton in the nucleus of an atom

Mass number

number of neutrons and protons in the nucleus

Isotope

atoms of the same element, with the same number of protons but different number of neutrons

Relative atomic mass

The average mass of an atom of an element, taking into account the abundance of all its isotopes. With reference to 1/12th the mass of Carbon-12

Relative Atomic Mass formula

(% of isotope A x mass of isotope A) + (% of isotope B x mass of isotope B) / 100

Sub-atomic Particle Positions

• Protons and Neutrons: in nucleus

• Electrons: in shells

Sub-atomic Particles Relative Charges

• Protons: +1

• Electrons: -1

• Neutrons: 0

Element

• a substance made of atoms that all contain the same number of protons and can’t be split into anything simpler

• there are 118 elements

• ex: Copper, Iron, Arsenic

Atom

Smallest part of an element that has the element's properties.

Compound

• a pure substance made of two or more different elements chemically bonded

• ∞ number of compounds

• can’t be separated by physical means

• ex: NaCl, MgO

Mixture

• combination of two or more substances mixed but not chemically bonded together.

• can be separated by physical means like filtration/evaporation

• ex: sand and water, oil and water

Molecule

One or more element chemically combined

Heating Curve

Graph showing how a substance changes state when heated

Cooling Curve

A graph showing how a substance changes state when cooled.

Filtration

Used to separate insoluble solid impurities from a mixture.

Fractional Distillation: Use

• Used to separate miscible substances with different boiling points

• Can separate more than 2 substances

Miscible Liquids

• Liquids that mix

• They form one layer

• Eg: ethanol and water

Immiscible Liquids

• Liquids that don't mix

• They form more than one layer

• Eg: Oil and water

Properties of the compound vs Properties of the element it's made of

Different

Pure salt from Rock Salt

• Filtration and Crystallisation

• Rock salt = Pure salt + insoluble impurities

• Filtration: separates the insoluble impurities from the mixture

Element, Compound, Mixture Diagram

Pure Substance vs Mixture

• Natural language: natural and clean, nothing added

• Chemistry: pure substance = single element/compound with no other substances

• ex: pure water has only H₂O molecules

• drinking water isn’t pure because it has additional substances like dissolved ions and chlorine

Distinguishing Purity

• Pure substances melt and boil at specific temperatures                                              ex, pure water m.p. = 0^°C, b.p = 100°C

• impure substances have a range of melting and boiling points because they consist of different substances

• so, this data can be used to distinguish pure and impure substances

• this also helps us assess the purity of drugs and foodstuffs

Assessing Purity: Process

• using a melting point apparatus to heat a small portion of the sample and observe the exact melting point

• compare to a data table

• the closer the value is to the actual melting point, the purer the substance

Simple Distillation: Process

• separate a liquid and a soluble solid from a solution or a pure liquid from a mixture of liquids

• the solution is heated, and the liquid evaporates

• the vapour rises through the neck of the round-bottomed flask

• the vapour passes through the condenser, where it condenses and cools to a pure liquid that is collected in a beaker

Fractional Distillation

• used to separate two or more miscible liquids

• solution is heated to the temperature of the substance with the lowest m.p

• this substance will evaporate first, and vapours will pass through a condenser

• the liquid will be condensed and collected in a beaker

• once all of the substance is collected, the other liquid will be left behind

Fractional Distillation of Ethanol and Water

• Ethanol b.p = 78 °C, Water b.p = 100 °C

• mixture is heated until it reaches 78 c, and the ethanol boils and distills out of the mixture and condenses into the beaker

• when the temperature starts to increase to 100 °C, heating should be stopped. Water and ethanol are now separated

Filtration: Process

• used to separate undissolved solids from a mixture of solids and a liquid (centrifugation can also be used for this mixture)

• filter paper is placed in a filter funnel above a beaker

• the mixture is poured in

• the filter paper will only allow small liquid particles to pass through as filtrate

• solid particles are too large, so the stay behind as filtrate

Crystalllisation

• used to separate a dissolved solid from a solution, when the solute is more soluble in hot solvent than cold

• solution is heated, allowing the solvent to evaporate and leave a saturated solution behind

• test if the solution is saturated by dipping a glass rod into the solution (if the solution is saturated, crystals will form on the glass rod)

• saturated solution will cool slowly

• crystals will grow as solids come out of the solution due to decreasing solubility

• crystals are collected by filtering, they are washed with cold distilled water to remove impurities and then allowed to dry 

Paper Chromatography: Process

• used to separate substances with different solubilities in a given solvent

• a pencil line drawn on chromatography paper, and spots of the sample are placed on it (pencil is used as ink would run into the chromatogram along with the samples)

• the paper is then lowered into the solvent container (the line must be above solvent line so samples don’t wash into the solvent container)

• solvent travels up the paper by capillary action, taking the coloured substances with it

• different substances have different solubilities, so will travel at different rates    (this causes the substances to separate, higher solubility = travel further)

• this will show the different components of the ink/dye

Interpreting Chromatograms

• pure substances = only one spot

• impure substance = multiple spots

• same substance = identical chromatograms

• mixture = separates to show different components as separate spots

• use a known compound to identify spots

Rf Values

• used to identify components of mixtures

• always the same for a particular compound

• solvent changed = Rf value changes

• allows to identify compounds because it can be compared with known values

• Rf = distance moved by substance / distance moved by solvent

Law of Conservation of Mass

Total mass of reactants is always equal to total mass of products

RFM Calculation

• multiply each A_r by how much there is

• add them together

Avogadro’s Number

6.023 × 10²³

Moles and Mass

n = M/M_r

Calculating Molar Mass/ no. of moles

• calculate M_r

• divide mass given by M_r

Calculating Reacting Mass

• find reacting moles (given or calculate with M / M_r)

• find ratio between given substance and substance to find

• find no. of moles

• use no. of moles to find mass

Balancing Equations using Reacting Mass

• write unbalanced equation

• write down masses

• calculate moles using mass and M_r

• use moles to find ratio

• use ratio to balance equation

Reasons for not getting 100% yield in a reaction

• some reactants left behind

• reaction may be reversible- high yield is impossible since products are continually turning back into reactants

• some products lost during purification or separation stages like filtration or distillation

• side reactions: substances reacting with gas in the air or impurity in reactant

• products lost during transfer between containers

Thoretical Yield

• amount of product that would be obtained under perfect practical and chemical conditions

• calculated from balanced equation and the reacting masses

Actual Yield

• recorded amount of product obtained

• always less than theoretical yield

Percentage Yield: Purpose

• compares actual yield to theoretical

• for economic reasons, objective of every company is to have yield % as high as possible → to reduce costs & wastes and increase profits

• good way to measure how successful a chemical process is

Calculating Percentage Yield

• find actual yield (usually given)

• find theoretical yield (calculate using moles and mass)

• use equation: (actual/theoretical) x 100

Experiment (Finding Formulae of Simple Compounds): Aim

to determine formula of hydrated copper sulfate: CuSO₄.xH₂O

Experiment (Finding Formulae of Simple Compounds): Method

• measure mass of evaporating dish

• add known mass of hydrated salt

• heat over bunsen burner, gently stirring

• stop when salt turns from blue to white (all water lost)

• record mass of dish and contents

Experiment (Finding Formulae of Simple Compounds): Overheating the salt

decomposes and gives a larger change in mass

Experiment (Finding Formulae of Simple Compounds): Results

mass of white anhydrous salt

• measure mass of white anhydrous salt (mass of salt remaining)

mass of water

• subtract mass of white anhydrous salt from mass of known hydrated salt

• divide mass of the salt and water by relative masses (find moles)

• simplify the ratio (multiply by 2 if decimal)

• find ratio as 1:water

• represent ratio as ‘salt.xH₂O)

Practical (Determining Formula of Magnesium Oxide): Aim

To determine the empirical formula of magnesium oxide by combustion of magnesium

Practical (Determining Formula of Magnesium Oxide): Diagram

Practical (Determining Formula of Magnesium Oxide): Method

• measure mass of crucible with lid

• add Mg sample to crucible and measure mass with lid

• find mass of Mg

• strongly heat crucible over Bunsen burner for several minutes

• lift lid frequently to allow sufficient air into crucible for Mg to oxidise without letting MgO smoke escape

• continue heating until mass remains constant (max mass) → reaction is complete

• measure mass of crucible and contents

• calculate mass of crucible and contents by subtracting mass of empty crucible

Practical (Determining Formula of Magnesium Oxide): Results

• find mass of metal by subtracting mass of crucible from Mg and mass of empty crucible

• subtract mass of Mg used from mass of MgO

• divide each mass by A_r,

• simplify ratio (multiply by 2 if decimal)

• represent as M_xO_y

Practical (Determining Formula of Copper(II) Oxide): Aim

To determine the formula of copper(II)oxide by reduction with methane

Practical (Determining Formula of Copper(II) Oxide): Diagram

Practical (Determining Formula of Copper(II) Oxide): Method

• measure mass of empty boiling tube

• place metal oxide into a horizontal boiling tube and measure mass again

• support tube in horizontal position by clamp

• natural gas(methane) is passed over copper(II) oxide and excess gas is burned off

• copper(II) oxide is heated strongly with a Bunsen burner

• heat until metal oxide fully changes colour (all oxygen removed)

• measure mass of remaining powder in the tube and subtract mass of tube

Practical (Determining Formula of Copper(II) Oxide): Results (Empirical Formula)

• measure mass of powder to find mass of metal

• divide masses by A_r

• simplify ratio

• represent ratio as M_xO_y

Molecular Formula: Definition

Formula showing number and type of each atom in a molecule. Ex: ethanoic acid is C₂H₄O₂

Empirical Formula: Definition

Simplest whole number ratio of atoms of each element present in one molecule or formula unit of the compound. Ex: ethanoic acid is CH₂O

Ionic compounds are always _________ formulae

Empirical

Calculating Empirical Formulae

• write element

• write value given (% or mass)

• write A_r

• calculate moles by m/M_r

• calculate ratio of moles (multiply to make all values whole numbers)

• write final empirical formula

Calculating Molecular Formula

• find M_r of empirical formula (add masses of all atoms in the empirical formula

• divide M_r of molecular formula by M_r of empirical formula

• multiply each number in empirical formula by answer to find molecular formula

Calculating Concentration of Solutions in mol/dm³

number of moles (mol) / volume (dm³)

Avogadro’s Law

At the same conditions of temperature and pressure, equal amounts of gases will occupy the same volume of space

Molar Gas Volume at RTP

24dm³ or 24000 cm³

RTP

room temperature and pressure (20^oC and 1atm)

Gas Volume Equations

• volume = moles x 24 (dm³/mol)

• volume = moles x 24000 (cm³/mol)

Metals + Cold Water Reaction Speeds

K : violent

Na: quick

Li and Ca: less strong

Fe: slow rust

Mg, Zn, Cu: no reaction/very slow

Metal + Water Reaction Format

metal + water —→ hydroxide + hydrogen gas

Metals and Acids Reaction Rate

• only metals above Hydrogen

• more reactive metal = more vigourous reaction

• K and Na are very dangerous and react explosively

Metal + Acid Reaction Format

Metal + Acid → Salt + Hydrogen

Solid

• regular arrangement

• low kinetic energy

• vibrate in fixed positions

• strong forces of attraction

Liquid

• arranged close together

• more kinetic energy than solid particles

• move past each other

• weak forces of attraction

Gas

• arranged far apart

• a lot of kinetic energy

• move randomly in all directions

• very weak forces of attraction

Evaporation

• liquid turns into a gas

• occurs at any temperature

• faster, because it is a surface level process

Boiling

• liquid turns into a gas

• only at a certain temperature

• slower, because all the particles need to overcome forces of attraction

Diffusion

Spreading out of particles from high concentration to lower concentration until there is equal concentration throughout.

Diffusion: temperature effect

• Increases with increase in temperature

• Because when particles are heated, they gain kinetic energy and move around more

Bromine Experiment (and Hydrogen variant)

• Lower jar: Bromine gas

• Top jar: air

• When lids are removed, bromine diffuses upwards

• Air also diffuses downwards until both jars are uniformly brown

• Hydrogen variant

  • place a lighted splint to check

  • Expected: sound from top jar because H is less dense

  • But equal sounds from both jars

Melting

• solid to liquid

• heat energy absorbed transformed to kinetic energy

• happens at a specific temperature or melting point (m.p.)

Freezing

• liquid to solid

• reverse of melting, happens at (m.p.)

• needs a significant decrease in temperature and loss of thermal energy

Sublimation

• solid to gas directly

• very few solids, eg: iodine, carbon dioxide

• reverse reaction: desublimation, deposition

Potassium Manganate (VII) Experiment

Description

• when crystals are dissolved, a purple solution is formed

• a small number of crystals produce a highly intense colour

Explanation

• water and potassium manganate (VII) particles move randomly and slide across each other

• therefore mix easily

• diffusion in liquids is slower than in gases because particles are closer and move more slowly

NH₃ + HCl Reaction

• particles diffuse along the tube

• white ring of NH₄Cl forms where they meet

• ring forms closer to HCl end

• because Ammonia particles are lighter and move faster

• so they travel further in the same amount of time, and further away from the NH₃ end

Dilution of Potassium Magnate (VII)

Description

• solution can be diluted several times

• colour fades but doesn’t disappear until after many dilutions

Explanation

• indicates that there a lot of particles in a small amount of potassium magnate (VII) and so the particles are very small

Solvent

• liquid in which a solute dissolves

• eg: water in salt water

Solute

• substance that dissolves in a solvent to become a solution

• eg: salt in salt water

Solution

• mixture formed when a solute is dissolved in a solvent

• eg: sea water

Saturated solution

• a solution with the max. conc. of solute dissolved in the solvent

• eg: sea water in the Dead Sea

Soluble

• describes a substance that will dissolve

• eg: salt is soluble in water

Insoluble

• describes a substance that will not dissolve

• eg: sand is insoluble in water

Solubility

measurement of how much of a substance will dissolve in a given volume of a liquid

Solubility of Solids & Temperature

temperature increase = solubility increases

Solubility of Gases (Temperature and Pressure)

• pressure increase = more soluble

• temperature increases = less soluble

Solubility Practical

• pour tap water into a beaker

• heat to a specific temperature

• add solute one spatula at a time, with constant stirring until no more dissolves and some remains undissolved in the mixture

• monitor temperature and keep it uniform throughout

• record mass of empty evaporating basin

• filter the mixture of of solution and undissolved crystals into the evaporating basin

• evaporate the filtrate until dry crystals are formed

• record the mass of the evaporating basin with the dry crystals

Groups

• vertical columns

• number of electrons in the outermost shell

• 1-7(0)

Period

• horizontal rows

• show the number of shells

• 1-7

Electronic Configuration in Shells

• first shell: 2

• second shell: 8

• third shell: 8