3.1 & 2 Flashcards
How is the periodic table arranged
By increasing atomic no.
In periods showing repeating trends in physical and chemical properties
In groups having similar chemical properties
First ionisation energy
Removal of 1 mole of electrons from each atom in 1 mole of gaseous atoms
Metallic bonding
Strong electrostatic attraction between cations and delocalised electrons
Solid giant covalent lattices
Networks of atoms bonded by strong covalent bonds
Atomic radii across a period
Decreases
More attraction to nucleus
Greater no. of protons
When is the ionisation energy higher
With a smaller radius (higher nuclear charge)
Greater attraction
More energy needed
How do inner quantum shells affect the IE
Inner quantum shells shield the electrons from the attractive force of the nucleus
The more inner quantum shells
The greater the shielding effect
The lower the IE
When a molecule sublimates what bonds are broken
The intermolecular forces not the covalent bonds
Why is Mg harder than Na
No. of delocalised electrons
Charge on cations
Cations are smaller, so the charge density is even greater
Why is there an increase in mp and bp of metals from Gp 1 to Gp3
Increased strength of metallic bonding: The charge on the positive ion is greater More delocalised ions (1 in 1+ vs 3 in 3+) Smaller ions (greater charge density)
Properties of metals
Malleable
Ductile
High mp and bp
Good electrical conductors
Solubility of giant ionic lattices
Ionic lattices dissolve in polar solvents like water
Cations attract delta negative charges on the O molecule in H2O
Anions attract delta positive charges on the H molecules in H2O
Solubility of simple molecular lattices
London forces can form between molecules and non-polar solvents (hexane), weakening the lattice
Group 2 characteristics
Fairly high melting and boiling points
Low densities
Group 2 reactions
2 Mg (s) + O2 (g) —> 2MgO (s) Ca (s) + 2 H20 (l) —> Ca(OH)2 + H2
Second IE
X (g) —> X^+ (g) and e^-
X^+ (g) —> X^2+ and e^-
Why does reactivity increase down Group 2
Further from nucleus
More shielding
1st and 2nd IE decreases
Why does solubility and pH increase down a group
Mg(OH)2 —> Mg2+ and 2 OH-
Dissociation causes more hydroxide ions to be released —> more alkaline solution
Down Group 7
No. of full shells increase
Outermost electrons further from nucleus
Atomic radius increases
Forces in Group 7 molecules
London forces
More electrons = stronger London forces
What do halogens react to form
Halide ions
Are halogens reduced or oxidised
Reduced
Oxidising agent
Why does reactivity decrease down Group 7
Nuclear charge increases down a group
Atomic radius increases
Shielding increases (outweighs nuclear atraction)
Attraction between nucleus and electron decreases
Halogen less readily accepts an electron
Do halide ions get oxidised or reduced
Oxidised
Why is iodide the most powerful reducing agent
It wants to lose electrons
How does chlorine react with water
To form an acidic solution containing HCl and HOCl (chloric (I) acid)
Risks of chlorinating water
Cl2 gas is toxic
Chlorinated hydrocarbons may form (carcinogens)
Benefits of chlorinating water
Effectively kills many harmful microorganisms
Reduction in cholera and typhoid rates
Bleach
Formed when Cl2 is added to cold dilute NaOH solution
Active ingredient is sodium chlorate (NaOCl)
Cl2 (aq) + 2 NaOH (aq) —> NaOCl (aq) + Nacl (aq) + H2O (l)
Tests for halide ions
AgNO3 used to test
Silver halides are insoluble
Cl- (aq) and Ag+ (aq) —>
AgCl (s)
White precipitate
Br- (aq) and Ag+ (aq) —>
AgBr
Cream precipitate
I- (aq) and Ag+ (aq) —>
AgI
Yellow precipitate
What does AgCl dissolve in
Dilute ammonia
AgBr dissolves in
Concentrated ammonia
Does AgI dissolve
No
Colours with cyclohexane
Cl2 - pale green
Br2 - orange
I2 - violet
Order of identifying ions
Carbonate
Sultate
Halide
Carbonate test
Dilute nitric acid
Bubbles of gas, effervescence if positive
If you bubble CO2 through lime water it goes from clear to cloudy
Sulfate test
Add Ba^2+ (Ba(NO3)2)
Forms white precipitate if positive - barium sulfate
Halide test
Dissolve halide in H2O and add AgNO3
AgX precipitate colour if positive
Also test if dissolves in NH3
Ammonium test
Add warm NaOH and place on damp litmus paper
Turns blue
Why is the order of the ion tests important
BaCO3 (s) and Ag2SO4 (s) are both insoluble white precipitates but only BaCO3 would give off bubbles of gas so you’d automatically know
If you have a mixture of ions:
Do tests in same order
Keep adding HNO3 until reaction stops
Add excess of barium nitrate to ensure all SO4^2- form a precipitate - which can be filtered
Then test filtrate for halide ions
Why is dilute nitric acid used in the carbonate test
It does not form any precipitates whereas HCl and H2SO4 do
Direct measurement of enthalpy change
Q = m * c * delta-t Delta-H = -Q/n
Use n of moles that is not in excess
Hess’ law
Enthalpy is a state function, depends only on state
If a reaction can take more than one route and the initial and final conditions are the same, the total enthalpy change is the same
Catalyst
A substance that increases the rate of a chemical reaction without being used up in the overall reaction.
Does this by providing an alternative route with a lower activation energy
Homogeneous catalysis
A reaction in which the catalyst and reactants are in the same physical state usually aq or g
Heterogeneous catalysis
A reaction in which the catalyst and the reactants are in diff. states. The reactant is usually a gas and the catalyst, a solid
Dynamic equilibrium
The equilibrium that exists in a closed system when the rate of the forward reaction is equal to the rate of the reverse reaction and the conc. don’t change
Le Chateliers principle
When a system in dynamic equilibrium is subjected to a external change, the system readjusts itself to minimise the effect of the change and to restore equilibrium
What is Ca(OH)2 used for
Neutralising acid soils
What are Mg(OH)2 and CaCO3 used as
‘Antacids’ in treating indigestion
Delta H
Reactants - products
-ve delta h
Exothermic
+ve delta h
Endothermic
Activation energy
Minimum energy required to start a reaction
Enthalpy change of formation
Formation of 1 mol of a compound from its constituent elements, in standard conditions
Products and reactants in their standard form
Standard enthalpy of neutralisation
Formation of 1 mol of water from neutralisation
Standard conditions
298K
100 kPa
Collision theory
For a reaction to take place, particles must collide in the correct orientation with sufficient energy
Why does using a catalyst increase the rate
It provides an alternative route with lower Ea so a greater proportion of collisions have an energy > Ea
Increasing the temperature increases the rate because
A greater proportion of molecules have an energy > Ea and there are more frequent successful collisions
Changing the pressure/conc. changes the rate because
There are more/less particles per unit volume and therefore there are more/less collisions
Enthalpy change of formation
Formation of 1 mol of a substance from its constituent elements under standard conditions and states
Enthalpy change of combustion
Complete combustion of 1 mole of a substance with excess O2 under standard conditions and states
Enthalpy change of neutralisation
Formation of 1 mol of water from neutralisation under standard conditions and states
How to measure rate of reaction from a graph
Drawing a tangent at said point and calculating the gradient
Factors affecting rate of reaction
Temperature
Concentration
Surface area
With or without catalysts
Effect of increasing surface area
All particles have exposed surfaces to react with each other; more frequent collisions
What does the highest point of the Boltzmann curve shows
Most probable energy
What is the total area under the Boltzmann distribution equal to
The no. of molecules
If the forward reaction is exothermic, how does increased temp affect the equilibrium
Shifts to the LHS
If the forward reaction is exothermic, how does decreased temp affect the equilibrium
Shifts it to RHS
How does increasing pressure affect the equilibrium
Shifts to side with least moles of gas
How does decreasing pressure affect equilibrium
Shifts to side with most moles of gas
How does the equilibrium shift if the conc of reactants increases
Shifts to the RHS
How does the equilibrium shift if the conc. of products increases
Shifts to the LHS
If the forward reaction is endothermic, how does increased temp affect the equilibrium
Shifts to the RHS
If the forward reaction is endothermic, how does decreased temp affect the equilibrium
Shifts to the LHS
What is Kc a measure of
The conc. of products and reactants at equilibrium
How is Kc calculated
Kc = [C]^c [D]^d/ [A]^a [B]^b
If Kc»_space; 1
More products are formed at equilibrium
If Kc «_space;1
More reactants formed at equilibrium
If Kc = 1
Equal conc. of reactants and products
Why do solid giant ionic structures not conduct electricity
They have ions fixed in position by ionic bonds and so don’t conduct
Role of a catalyst
Increasing rate of reaction without being used up by the overall reaction
Providing alternative route with lower Ea
Homogenous catalysts
Has same physical state as the reactants. Reacts with reactants to form intermediate. This then breaks down to give the product and regenerates the catalyst
Heterogenous catalyst
Has diff. physical state to reactants. Reactant molecules are absorbed onto the surface of the catalyst where the reaction take place, activation energy is lowered and the product molecules leave by desorption
Catalyst for Haber process
Solid Iron
Catalyst for reforming
Pt (s) or Rh (s)
Catalyst in hydrogenation of alkenes
Solid Ni
Industrial preparation of ethanol
C6H12O6 —> 2 C5H5OH + 2 CO2
Benefits of catalysts
Increased sustainability:
Lowering temps and reducing energy demand from combustion of fossil fuels with resulting reduction in CO2 emissions
Thermal decomposition of Group 2 carbonate
CaCO3 —> CaO + CO2
Periodicity
Repeating patterns of properties across diff. periods
Why do solid ionic lattices not conduct electricity but molten ones do
The ions are fixed in place by ionic bonds when solid but when molten the ions are mobile
Which Group 2 carbonate decomposes at the highest temp
BaCO3
Are giant metallic structures soluble in water
No
Bonding in gaseous hydrogen halides
Mainly covalent with an increasing tendency towards ionic as you go up the group
Which particles are attracted in metallic bonding
Cations and delocalised electrons
Which halogen most readily forms 1- ions
Fluorine
Percentage uncertainty
No. of readings * uncertainty/ (quantity measured) * 100
Limitation in enthalpy pag
Fuel evaporates
Heat loss to surroundings
Incomplete combustion.
Sources of error in titrations
Not accurately weighing mass of solid - use more precise balance
Not all acid gets transferred to volumetric flask - rinse out beaker and add rinsings to volumetric flask
Insufficient mixing of solution - invert several times
Burette not rinsed - rinse with acid solution before use
Which halogen most readily forms 1- ions
Fluorine
NO and CO in a catalytic converters
2NO + 2CO —> 2CO2 + N2
CO and NO adsorbed onto surface
CO2 and N2 made, lowered activation energy
Products desorb from the surface
General trend in mp and bp from Gp 1 to Gp 4
Increases
Why is there an increase in mp and bp between Gp 3 - Gp 4
From Al to Si; giant metallic lattice to giant covalent lattice
Covalent bonds are stronger than metallic bonds in the same period
Why is there a decrease in mp and bp between Gp 4 to Gp 7
Go from giant covalent lattice to simple molecular lattice w/ weak IMF (LF) and molecules get smaller e.g. P4 and Cl2
Simple molecular lattice
Weak IMF (LF) between molecules
The bigger the molecules …
The more electrons
Stronger induced dipole-dipole forces
Higher mp
How does repulsion between electrons in orbitals affect the IE
Less repulsion leads to larger IE
Why doesn’t a catalyst change the position of eqm
Increases the forward and backward reaction by the same amount
Enthalpy change of reaction
No. of moles of reactants specified in the eqn react together
Avg. bond enthalpy
One mole of bonds breaking
Why may the actual bond enthalpy be diff to the calculated value
Bonds have diff strengths in diff environments
Why do industrial manufacturing processes use catalysts
Higher atom economy Reduce CO2 emission Enable reactions to occur w/ more specificity (correct stereoisomer) Reactions can occur at lower temp Saves energy costs
What makes a reaction exothermic
Bond breaking absorbs energy
Bond forming releases energy
More energy released than absorbed
Endothermic
Energy enters system from surrounding
Energy profile diagram for endothermic reaction
Products above reactants
Ea and delta H going up to products
How is a dynamic eqm reached
Rate of forward reaction slows down and rate of backwards reaction speeds up
Ionic eqn of carbonate test
CO3 2- + 2H+ —> CO2 + H2O
Directly measuring enthalpy change of solution/ neutralisation
Measure out sol using a gradated pippette into a polystyrene cup
Weigh out excess of solid
Place thermometer into cup; stir and record temp every 30s for X mins
At X+1 mins add solid to cup
Gently stir and record temp every 30s until 10mins
Draw cooling curve and extrapolate
Why must we extrapolate cooling curves to get an accurate value for delta T
Reaction isn’t instantaneous
Heat is lost as mixture heats up
Why may enthalpy change of neut/sol be diff from accepted value
No lid is used - heat loss
Spp heat capacity may be inaccurate
Energy absorbed by polystyrene
