t1: sub-atomic particles atoms are made from smaller particles called subatomic particles. there are...
TRANSCRIPT
T1: Sub-atomic particles• Atoms are made from smaller particles called
subatomic particles.• There are three types we need to know about,
summarised below.
T1: Mendeleev• Arranged elements by increasing atomic mass
but….• He broke this rule and left some gaps if an
element’s properties weren’t similar to the one above it.
• He thought the gaps were for elements that hadn’t been discovered yet and predicted their properties.
• When they were discovered, the properties matched the predictions
PERIODS….increasing atomic mass, differing properties
GRO
UPS…
…sim
ilar properties
Element Type
= non-metal = metal
Particle Relative charge
Relative mass
Found?
Proton 1 Positive, +1 In nucleusNeutron 1 Neutral, 0 In nucleusElectron Neglible () Negative, -1 In shells orbiting
nucleus
T1: Reading the Periodic Table
• Note: on some periodic tables, they are the wrong way up, just remember that the smaller number is the proton number.
Relative Atomic Mass (aka nucleon number):
The total number of protons and neutrons added together.
Atomic number (aka proton number):
The number of protons or electrons.
T1: What’s in my atom?
Protons = atomic numberElectrons = atomic numberNeutrons = relative atomic mass . – atomic number
Atomic number = 9Relative Atomic mass = 19Protons = 9Electrons = 9Neutrons = 19-9 = 10
T1: Atoms and Elements•Element = substance containing only one type of atom.•Protons and electrons: same for every atom of an element…it is the number of protons that decides the element.•Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes
T1: Relative Atomic Mass•This is the mass of an element relative to 1/12th the mass of 12C.•Element: substance containing only one type of atom.•Protons and electrons: same for every atom of an element…it is the number of protons that decides the element.•Neutrons: can differ…atoms with the same number of protons but different numbers of neutrons are called isotopes.
T1: Isotopes (HT)• Versions of an element with same atomic number
but different atomic mass.• Number of protons is the same, but number of
neutrons is different.• Relative Atomic Mass is average of the masses of the
isotopes, weighted by their relative abundance
• For example, Neon has three isotopes
• Relative atomic mass of Neon =
• This is why some atoms have a relative atomic mass with a decimal point.
T1: Electron Configuration• Electrons orbit the nucleus in shells.• First shell holds two electrons• Second and third shell hold 8 electrons• Note: the third shell can actually hold more,
but we won’t worry about this until A-level.
Example: SiliconAtomic number is 14, so it has 14 electrons.You build up electrons from the first shell outwards, so in this case: - First shell has 2 - Second shell has 8 - Third shell has 4
This can be written as: 2.8.4; or drawn as:
Neon Isotope Mass
Relative Abundance (%)
20 90.5
21 0.3
22 9.2
Note: Si is in period three and group four of the periodic table; it also has three electron shells and four electrons in the outer shell – this is no coincidence!
T2: Forming IonsCations are positive (cat…pussitive!) ionsThey are formed when atoms lose electrons.Metals form cations by losing the electrons in their outer shellsIn the example, aluminium loses its three outer-shell electrons to become Al3+…each lost electrons cause 1 ‘+’ charge.
Anions are negative ionsThey are formed when atoms gain electrons.Non-metals form anions by filling their outer shells.Name ends with ‘-ide’ to show it is a negative ion,In the example, oxygen gains two outer-shell electrons to become O2-, giving it 8 electrons in its outer shell.
3+
T2: Making Ionic Compounds• An ionic bond is the attraction between a
positive and a negative ion.• The overall number of positive and
negative charges must cancel out.• Form between a metal and a non-metal• Ionic compounds do not form molecules
Example 1: Magnesium reacting with chlorine.• Anion: Cl forms Cl- ions• Cation: Mg forms Mg2+ ions• Formula = MgCl2 • Why: two Cl- gives a 2- charge to balance
2+ from Mg2+. • Name: magnesium chloride
Example 2: aluminium reacting with oxygen.• Anion: O forms O2- ions• Cation: Al forms Al3+ ions• Formula = Al2O3 • Why: Two Al3+ gives a 6+ charge, three
O2- gives a 6- charge. • Name: aluminium oxide
T2: Ionic Structures (HT)•A repeating 3D lattice of positive and negative ions.•Strong electrostatic bonds between ions.
T2: Precipitates and Precipitation•When an insoluble salt is formed from the reaction of two soluble salts.•Goes cloudy as small particles of solid are made.•Predicting precipitates: simply choose a combination of soluble salts where you tell that if the ions swapped over you would get an insoluble salt: use the solubility table for help.•Example:
T2: Common Ions• You should try to memorise the ions formed
by various species:
• There are also some ‘compound’ ions made of more than one atom with an overall charge:
• Hydroxide: OH-
• Nitrate: NO3-
• Sulphate, SO42-
• Carbonate, CO32-
• Ammonium, NH4+
Group Electrons in outer shell
Ion formed
Examples
1 1 + Li+, Na+, K+
2 2 2+ Be2+, Mg2+, Ca2+
3 6 2- O2-, S2-
4 7 - F-, Cl-, Br-, I-
T2: Solubility• Soluble: a compound dissolves in a given liquid.• Insoluble: a compound does not dissolve.
Soluble in water In soluble in waterAll sodium, potassium, ammonium saltsAll nitratesMost chlorides Except: silver and lead
chloridesMost sulfates Except: lead, barium
and calcium sulfates. Except: sodium, potassium and ammonium carbonates
Most carbonates
Except: sodium, potassium and ammonium hydroxides
Most hydroxides
T2: Properties of Ionic Compounds• Melting point: High due to strong bonds between ions.• Boiling point: Higher, due to strong bond between ions.• Solid: do not conduct electricity• Molten (liquid): do conduct electricity• Dissolved (aqueous): do conduct electricity
Why? (HT)Electrical Conductivity• Electricity is conducted when there are charged particles
that are free to move.• Solid: there are charged particles (the ions), but they are
not free to move, so they do not conduct.• Liquid/Aqueous: the ions are now free to move, so they
do conduct
High Melting/Boiling Points• Ionic bonds (attraction between positive and negative
ions) are very strong.• Melting and boiling require these bonds to be broken.• This takes lots of (heat) energy.
T2: Making Insoluble Salts
1. React solutions of (the right) two soluble salts together.
2. Filter the mixture to collect the precipitate.
3. Rinse the filter residue with distilled water to remove impurities.
4. Allow the residue to dry.
T2: Barium Meals
• A patient is given a drink containing barium sulfate.• This can show up on
a x-ray, helping doctors to investigate the digestive system.
T2: Flame tests1. Clean a metal loop in acid2. Did loop in a metal salt.3. Heat in roaring Bunsen flame.
• Sodium, Na+ Yellow• Potassium, K+ Lilac• Calcium, Ca2+ Red• Copper, Cu2+ Green-blue
Precipitation TestsChloride: add acidified silver nitrate to get a white precipitate if chloride is present.Sulfate: add acidified barium chloride to get a white precipitate if sulfate is present.
Carbonate Test1. Add acid to the sample2. Pass any gas produced through
limewater: will go cloudy if the sample contained carbonate
T3: Diamond vs Graphite (HT)Diamond:• Very hard, as all carbon atoms joined
with strong covalent bonds.• Used to make cutting tools• Insulator as all electrons locked-tight
in bonds, so can’t move.
Graphite:• Layers of hexagonal carbon mesh that
rub away from each other, as there are only weak forces between the layers.
• Used as a lubricant.• Conductor as the electrons between
the layers are free to move. This is very rare for a giant covalent structure.
T3: Separating Immiscible Liquids• Immiscible = when liquids do not dissolve in
each other….like oil and water, one floats on top of the other.• Can be separated with a separating funnel;
the denser layer is tapped-off at the bottom.
T3: Covalent Bonds• Form when non-metals share electrons
between them.• Attraction between each atom and the
shared electron pair.• Atoms share electrons to complete their
outer shells• One bond is formed for each ‘gap’ in the
outer shell• Bonding represented with dot-and-cross
diagrams showing only the outer-shell electrons.
Example 1: WaterEach hydrogen needs one more electron to complete it’s outer shell and the oxygen needs two more. Oxygen forms two single bonds: one to each hydrogen.
Example 2: Carbon dioxide (HT only)Carbon needs two more electrons to complete it’s outer shell and each oxygen needs two more. Carbon forms two double bonds: one to each oxygen.
H HO
O OC
T3: Covalent StructuresSimple Covalent Molecules• Molecule = A particle made of a
small group of atoms, covalently bonded together.
• Low melting and boiling point, due to weak attractive forces between molecules..
• Electrical insulator as no electrons free to move.
• Examples: water, ammonia, oxygen
Giant Covalent• Repeating pattern of
many millions of atoms covalently bonded.• High melting/boiling
point because much heat energy needed to break strong covalent bonds.• Electrical insulator as
no electrons free to move.• Examples: silicon
dioxide, diamond, graphite
Lead nitrate + potassium iodide lead iodide + potassium nitratePb(NO3)2(aq) + 2KI(aq) PbI2(s) + 2KNO3(aq)
T3: Separating Miscible Liquids•Miscible = when liquids dissolve in each other…like alcohol and water.•Separate with fractional distillation using a fractionating column.•The components of the mixture have different boiling points, so if you heat it, each component will boil at a different time, allowing you to collect and condense the pure vapour.•We can do this to separate the gases in air by first cooling the air to turn the gases to liquid.
.
T4: Halogens and Their Reactions• Group 7: Fluorine (F) – pale yellow gas, Chlorine (Cl) – pale
green gas, Bromine (Br) – orangey-brown liquid, Iodine (I) – grey solid.• Most reactive at top of group, and get less reactive as you go
down.• Form halide ions with a charge of ‘-1’
Reaction with metals• React with metals to form metal halides• General equation: metal + halogen metal halide• For example: magnesium + iodine magnesium iodide Mg(s) + I2(s) MgI2(s) Note: Mg forms a 2+ ion, so two I- ions are needed.
Reaction with hydrogen• React with hydrogen to form hydrogen halides.• Hydrogen halides dissolve in water to form acids.• General equation: metal + halogen hydrogen halide• For example: hydrogen + fluorine hydrogen fluoride H2(g) + F2(g) 2HF(g) Note: hydrogen fluoride dissolves to make hydrofluoric acid.
Displacement Reactions• More reactive halogens can react with the ions of less
reactive halogens and displace them from compounds.• For example: 2KI(aq) + Br2(aq) 2KBr(aq) + I2(aq)
• This reaction works because bromine is more reactive than iodine.
• The orange colour of bromine would change to the brown colour of aqueous iodine.
• The reverse reaction would not work.
Reactivity Series of Halogens• Displacement reactions can be used to determine the
order of reactivity of the halogens.• Try reacting each halogen with solutions of each halide
salt, the halogen that does most reactions is most reactive.
T4: Metallic Bonding• Electrons are delocalised, moving
freely between all the atoms creating a ‘sea of electrons’• All atoms have a positive charge
as their outer-shell electrons have left them.• The bond is the attraction
between the positive ions and the sea of electrons.• Conduct electricity as electrons
are free to move.• Malleable (change shape but
don’t shatter when hit) because rows of atoms slide past each other when hit
T4: Transition Metals
• High melting points• Form brightly coloured
compounds
Halide SaltPotassium
fluoridePotassium chloride
Potassium bromide
Potassium iodide
Halogen
Fluorine x Reaction Reaction ReactionChlorine No
reactionx Reaction Reaction
Bromine No reaction
No reaction x Reaction
Iodine No reaction
No reaction No reaction
x
T4: Alkali Metals
• Group 1: Lithium (Li), Sodium (Na), Potassium (K)…• Properties: low melting point, soft (can be cut with
a knife).• React with water as follows:
General equation: metal + water metal hydroxide + hydrogenFor example: 2K(s) + 2H2O(l) 2KOH(aq) + H2(g)
Reactivity• Reactivity increases down the group:• Lithium just fizzes before disappearing• Sodium fizzes and gets hot enough to melt
into a ball, occasionally catching fire• Potassium fizzes very vigorously, getting hot
enough to burn with a lilac flame
Explaining Reactivity (HT only)• All reactions require you to remove the outer-shell
electron/• Atoms get bigger going down the group outer-
shell electrons further from nucleus easier to remove the outer shell electron.
T4: Noble Gases• Group 0 in the periodic table.• Helium ((He, Neon (Ne), Argon (Ar), Krypton (Kr)
Xenon (Xe), Radon (Rn)• Full outer shells so extremely unreactive: inert.
Discovery:• Lord Rayleigh noticed the density of nitrogen made in
reactions was less than nitrogen made from air.• Sir William Ramsey hypothesised that the nitrogen in
the air must also contain a denser gas that had not yet been discovered.• Through careful experiments, Rayleigh and Ramsey
discovered a gas that they named ‘argon’.• They also discovered helium, and then later Ne, Kr and
Xe.
Uses:• He and Ar were used to stop in filament in old bulbs
burning.• Ar and He used in welding to stop hot metal oxidising.• Ar used in fire extinguishing systems in server rooms.• He used in airships/blimps due to low density.• Neon lights due to red colour of light produce by neon.
T5: Endothermic and ExothermicExothermic Reactions• Chemical energy is converted to heat energy.• The surroundings get hotter.• For example: combustion reactions: Methane + oxygen carbon dioxide + water CH4 + 2O2 CO2 + 2H2O• Explosions are just very fast exothermic
reactions.
Endothermic Reactions• Heat energy is converted to chemical energy.• The surroundings get colder.• Examples: ammonium nitrate dissolving in
water, photosynthesis
Making and Breaking Chemical Bonds• In reactions, old chemical bonds are broken,
and then new ones are made.• Breaking bonds takes in energy; making bonds
gives out energy.• Stronger bonds take more energy to break,
and give out more when made.• In exothermic reactions, weaker bonds are
broken and stronger bonds are made.• In endothermic reactions, stronger bonds are
broken and weaker bonds are made.
Energy Diagrams (HT only)
Chem
ical
Ene
rgy
Reactants
Products
Energy released so gets hotter
Reactants
Products
Energy absorbed
so gets colder
EXOTHERMIC ENDOTHERMIC
T5: Catalytic Converters
• Part of exhaust pipe that helps make car exhaust less environmentally damaging.
• Toxic carbon monoxide and unburned hydrocarbons (from petrol) are converted into carbon monoxide and water.• The catalytic converter has a fine honeycomb
structure coated with the catalyst.• The catalyst contains a mixture of platinum,
rhodium and palladium.• The metals are expensive, so only a very thin
coating is used.• The catalysts work best at high temperatures, so
car exhaust is more damaging when the car has only just started and hasn’t warmed up.
T5: Collision Theory (HT)• To react: particles must collide with enough
energy.• To increase rate: increase the amount of
collisions or the energy of the collisions.
Effect of Concentration:• Increasing concentration increases the number of
reacting particles.• This increases the number of collisions.
Effect of Surface Area:• Increasing the surface area increases the
proportion of (solid) particles available to react.• This increases the number of collisions.
Effect of Temperature:• Increasing the temperature increases the speed
that particles are moving• This means there are more collisions, and those
collisions have more energy.
T5: Rates of Reaction (Intro)
Note: you increase the surface area by breaking a large piece into many smaller pieces, with powder being the best.
T6: Reacting Quantities (HT)• Combining relative masses with balanced equations lets us
work out the masses of chemicals involved in reactions.• We can use this mathematical relationship:
Example:• What mass of carbon dioxide can be produced by burning
15g ethene (C2H4) in excess oxygen (O2)?C2H4 + 3O2 2CO2 + 2H2O
• Substance 2 will be ethene, substance 1 will be carbon dioxide.• Calculate relative masses:
• Mr(ethene) = 2 x 12 + 4 x 1 = 28• Mr(carbon dioxide) = 12 + 2 x 16 = 44
• Then:
• m = mass of substance present• Mr = relative formula mass of
substance• n = number of substance in balanced
equation• 1 refers to the first substance• 2 refers to the second substance
Write out the equation.
Sub in the numbers
Rearrange to make m1 the subject.
T6: Empirical Formulae
Relative Atomic Mass, Ar• The lowest whole number ratio of atoms in a molecule.• For example:
• The empirical formula can be calculated from the masses of substances that react with each other as below.
• For example: 10.0g of magnesium reacts with 133.3 g of bromine.
Molecular Formula Empirical Formula
Water, H2O H2O
Ethane, C2H6 CH3
Glucose, C6H12O6 CH2O
Mg Br
Mass in g 10.0 133.3
Relative atomic mass 12 80
Divide by relative atomic mass
10 / 12 = 0.83 133.3 / 80 = 1.67
Divide both sides by smallest answer
0.83 / 0.83 = 1 1.67/0.83 = 2
Empirical formula MgBr2
T6: Yield• Theoretical yield: the amount of
product you would expect according to the calculation in the ‘Reacting Quantities’ box.• Actual yield: the amount of product
you actually get in practice.• Percentage yield: the proportion of
the theoretical yield that you actually achieve.
% yield is always less than 100 because:• The reaction may be incomplete• Some product may be lost during the
steps to prepare it.• Some reactants may also produce
products other than the desired one.
𝑚1
𝑀𝑟 1𝑛1=
𝑚2
𝑀 𝑟2𝑛2
T6: Percentage by Mass• This is the percentage of the mass of a compound due to a
particular element.
For example: what is the carbon in ethanol, C2H6O?
Calculate Mr of C2H6O Mr = (2 x 12) + (6 x 1) + 16 = 46
Number of C in C2H6O 2
Relative atomic mass of C
12
Percentage by mass of C = 52.1%
T6: Relative MassesElement Relative Mass
Hydrogen, H 1
Carbon, C 12
Oxygen, O 16
Sodium, Na 23
Chlorine, Cl 35.5
Relative Atomic Mass, Ar• The mass of atom relative to
the mass of 12C (carbon-12).• For example…
Relative Formula Mass, Mr• This is the sum of all the relative masses in a formula.• Relative formula mass of carbon dioxide, CO2:
Mr = Ar(C) + 2 x Ar(O) = 12 + (2 x 16) = 44
• Relative formula mass of sodium chlorate, NaClO3
Mr = Ar(Na) + Ar(Cl) + 3 x Ar(O) = 23 + 35 + (3 x 16)
• The rate of a reaction is its speed, how quickly products are made.• Reactions happen when particles collide with each
other.• Concentration: increasing concentration (the amount
of solute (dissolved stuff) in a given volume) will increase the rate.• Temperature: increasing temperature will increase
the rate.• Surface area: increasing surface area will increase the
rate.
Type of Bonding
Ionic Simple molecular Giant Molecular
How the bonds form
Swapping electrons to form ions
Sharing electrons Sharing electrons
Examples Sodium chloride, magnesium oxide
Water, methane, nitrogen
Quartz (silicon dioxide)
Bond strength
Strong Strong bonds, weak intermolecular forces
Strong bonds
Melting and boiling point
High Low High
Solubility Most in water Some in water Insoluble in water
Conduct electricity?
Only when molten or dissolved
No No (except graphite)