AS CHEMISTRY OCR A F321

ATOMS, BONDS AND GROUPS REVISION

Atomic Structure

The nucleus contains protons (positively charged) and neutrons (neutrally charged).

The atomic number (proton number) is equal to the number of protons in the atom’s nucleus and the amount of electrons orbiting the nucleus

The mass number is the total number of protons and neutrons in the nucleus.

Ions do not have the same number of electrons as protons, and so have an overall charge. Name Relative charge Relative massproton + 1 1neutron 0 1Electron -1 1/2000

Isotopes and Relative Masses

Isotopes are atoms that have the same number of protons but a different number of neutrons.

The relative atomic mass is the weighted mean mass of an atom of an element relative to C12.

The relative formula mass of a compound is equal to the sum of the individual relative atomic masses.

How to calculate an isotopes relative atomic mass…(Isotopes Mr x % abundance) +(the isotopes other Mr x % composition) 100

The Mole

a mole of a substance contains the same number of particles as there are in 12.0g of carbon-12 (6.023 x 1023

particles, also known as Avogadro’s number) One mole of a substance is simply the relative formula mass

for a compound, or relative atomic mass for an element in grams.

The empirical formula is the simplest whole-number ratio of atoms of each element present in a compound.

The molecular formula is the actual number of atoms of each element in a molecule.

Steps to calculate the empirical formula1. Find the moles of the known substance 2. Find the molar ratio 3. Find the moles of the other substance4. Change to its simplest form (empirical formula)

Molecular formula Molecular formula is the actual number of atoms of

each element in a molecule. You can determine the molecular formula of a molecule from

the empirical formula by… Diving the molecular mass of the molecule by the relative

molecular mass, then multiplying each molecule by this number.

Equations involving the Mole, concentration and volume

Moles= concentration x volume Mass (g) = Moles x Molar mass Volume = Concentration x 24dm3

Steps to calculate masses made in a reaction1. Balance equation2. Calculate moles of known3. Work out mole ratio of unknown 4. Convert moles to grams

Steps to calculate volume in dm 3 or cm 3 1. Divide mass by molar mass to give the amount of moles2. Multiply the amount of moles by 24dm3 , this gives volume

in dm3

3. But if you need to calculate volume in cm3 multiply by 1,000

Acids and Bases

An acid is a hydrogen ion (H+) or proton donor in solution,

A base is a hydrogen ion or proton acceptor in solution. Hydrochloric acid (HCl), sulfuric acid (H2SO4) and nitric acid

(HNO3) are common acids. Bases include metal oxides (e.g. MgO) metal hydroxides (e.g.

NaOH) and ammonia (NH3). Alkalis are soluble bases and form hydroxide ions OH-,

in solution. metal oxides, metal hydroxides, metal carbonates and

ammonia are all bases

To convert from dm3

to cm3

Multiply by 1,000!

Metal oxides contain the oxide ion (O2-). An oxide ion can accept two protons to form a water molecule: O2- + 2H+ H2O

Metal hydroxides contain the hydroxide ion (OH-). A hydroxide ion can accept a proton to form a water molecule: OH- + H+ H2O

Metal carbonates contain the carbonate ion (CO32-). A

carbonate ions can accept two protons to form carbon dioxide and water: CO3

2- + 2H+ CO2 + H2O Ammonia can accept a proton to form an ammonium ion: NH3

+ H+ NH4+

Reactions of Acids and Bases

The reaction between an acid and a base is called a neutralisation reaction.

Salts are formed when a hydrogen ion from the acid is replaced by a metal ion, or an ammonium ion.

Acids react with bases to form a salt and water only; they react with metal carbonates to form a salt, water and carbon dioxide gas.

Metals react with acids to form a salt and hydrogen gas. Salts may chemically combine with water as water of

crystallization in hydrated salts. Water of crystallisation can be removed by heating.

Reactants ProductsAcid + base Salt + waterAcid + Carbonate Salt + water + carbon dioxideAcid + alkali Salt + water

Observing neutralisation reactions

There is no visible change when an acid reacts with a soluble metal hydroxide or ammonia as all the reactants and products are colourless solutions. You need an indicator to determine that a reaction is taking place.

When an acid reacts with a solid metal hydroxide or oxide, the solid dissolves in the acid to form a solution.

When an acid reacts with a solid metal carbonate, the solid dissolves in the acid to form a solution and bubbles/fizzing/effervescence is seen

When an acid reacts with an aqueous metal carbonate, bubbles/fizzing/effervescence is seen

Salts and hydrated salts

A salt is an ionic compound that results from the neutralization reaction of an acid and a base, when the H+ ion of the acid is replaced by a metal ion or another positive ion such as ammonium (NH4+)

When an acid reacts with a base, it loses its H+ ion. The H ion is replaced either with the metal ion from the base or with an ammonium ion. The resulting compound is called a salt.

An anhydrous salt is a salt that contains no water. Hydrated salts become anhydrous salts once they are heated.

Salts are composed of anions (negative ions) and cations (positive ions)

Water of crystallisation occurs in salts, it is the water molecules that are trapped inside of hydrated salts

A hydrated salt is a salt that contains loosely bonded water molecules, this form when salts crystallise from a solution

Water of crystallisation are separated by a dot formulae.g. CuSO4.5H2O , This shows that for every mole of CuSO4

there are 5 moles of water. The number of moles of water of crystallisation is different

for every salt, but it must always be a whole number!

How to calculate the composition of hydrated salts1. Enter the % then divide by the molar mass to give moles2. Then divide the moles by the smallest to get the molar

ratio3. From this you can then derive the dot formula,which must

be in whole numbers!

Using mass data to find the value of X in hydrated salts 1. Enter the mass in grams 2. Then divide the mass by the molar mass to give the

amount of moles3. Then divide both of the moles by the smallest to get the

molar ratios4. From this you can then derive the dot formula, which must

be in whole numbers!

Acid-Base titrations These are carried out to discover how much acid is needed to

neutralise a base or vice versa Volume is measured in cm3

Concentration is measured in mol/dm3

How to carry out a titration1. Using a pipette add a measured volume of one

solution to a conical flask.2. The other solution is placed in a burette.3. The solution in the burette is added to the solution in

the conical flask, until the reaction has been completed. This is called the end point of the titration.

4. Then the volume of the solution added from the burette is measured.

The end point of the titration is identified using an indicator. Indicator Colour in

acidColour in base

End point colour if acid has been added to base

Methyl orange

red yellow orange

Bromothymol blue

yellow blue Green

phenolphthalein

colourless pink pink

To find out the concentration of the acid/base…1. Balance the equation2. Divide the volume by 1,000 to turn it from cm3 to dm3

3. Multiply the volume in dm3 by the concentration in mol/dm3 to find the moles of the known substance

4. Then you can find the molar ratio5. Multiply or divide by the molar ratio to find the

concentration of the unknown

Redox reactions A redox reaction is a reaction which involves oxidation

and reduction. Oxidation is the loss of electrons (in terms of electrons) Reduction is the gain of electrons (in terms of electrons) Oxidation is also the increase in oxidation number

during a reaction (in terms of oxidation number) Reduction is the decrease in oxidation number (in terms

of oxidation number) You can find out what element is being oxidised or reduced in

a reaction by finding out their oxidation number, before and after the reaction.

Oxidising agents cause oxidation by accepting an electron A reducing agent causes reduction by donating an electron You must be able to express half equations for redox

reactions e.g. during the reaction between Mg + Cl2 -> MgCl2

Mg is oxidised Mg-> Mg2+ + 2 e- And Cl is reduced Cl + e- -> Cl-

Oxidation Numbers

An oxidation number indicates the formal charge of a chemically combined particle in a compound.

The oxidation number is also the number of electrons that an atom uses to bond with atoms of another element.

The oxidation number of metals usually equals the group number (as a positive value) and minus (8-group number) for non-metals.

An element has been oxidised if the oxidation number increases, and reduced if the oxidation number decreases.

When they react, metals are normally oxidised (they lose electrons), whereas non-metals gain electrons and are reduced.

Oxidation number rulesAn element in its natural state and compounds have oxidation numbers of 0The oxidation number of a monatomic ion has the same charge as the ionHydrogen’s oxidation number is +1 (in a compound)Oxygen’s oxidation number is -2 (in a compound)Group 1 elements have an oxidation number of +1Group 2 elements have an oxidation number of +2Group 3 elements have an oxidation number of +3Group 7 elements have an oxidation number of -1The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

EXCEPTIONS TO THE RULE When bonded to fluorine, oxygen has an oxidation

number of +2. In peroxides, oxygen has an oxidation number of -1. (as fluorine is the most electronegative element)

When bonded to metals in hydrides, hydrogen has an oxidation number of -1.

Oxidation numbers in formulae Compounds have no overall charge so the sum of the

oxidation numbers must equal zero. Molecular ions have an overall charge, so the sum of the

oxidation numbers in the ion must equal the charge.

Oxyanions These are negative ions which contain an element along

with oxygen.

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Electronic Structure

Electrons occupy energy levels around the nucleus of the atom, where each shell has a principal quantum number.

For principal quantum number, n=1, the number of electrons is 2; for n=2, the number is 8; then 18; then 32 electrons for n=4.

Main energy levels are sub-divided into sub-shells and these consist of orbitals called s, p and d-orbitals.

Elements have an electronic configuration that can be shown in s, p or d notation, for example, sodium is 1s2,2s2 ,2p6 ,3s1 .

S orbitals are spherical and hold 2 electrons in opposite spins P orbitals are lobe shaped and hold 6 electrons in opposite

spins D orbitals hold 10 electrons F orbitals hold 14 electrons A way to remember electronic structure – Sophie, Simmonds’,

Penguin, Skips, Past, Sophie’s, Duck, Pedro

Energy level Number of electrons1st 22nd 83rd 184th 325th 50

Types of Bonds Ionic bonding takes place when positive ions and

negative ions are attracted in a giant ionic structure.

Xidation

(takes place between a non-metal and a metal ion, these ions have FULL charges)

Covalent bonding is the sharing of electron pair(s) between nuclei of atoms (takes place between non-metals and non-metals)

A polar covalent bond occurs when electrons are not shared equally between the atoms, as one has a greater electronegativity than the other.

A dative covalent bond is one formed in which both electrons are donated from the same atom.

A metallic bond is formed when positive metal cations are attracted to negatively charged delocalised electrons.

Delocalized electrons are electrons that are shared by TWO or more atoms.

Hydrogen bonding is the electromagnetic attractive interaction of a hydrogen atom and an electronegative atom, such as nitrogen, oxygen or fluorine.

Van der Waals forces are Weak electrostatic attractive forces between uncharged molecules, arising from the interaction of permanent dipoles

Ionisation energy Ionisation energy is the amount of energy needed to

remove one electron from each atom, in a mole of atoms, in the gas phase.

There are 3 factors that affect the amount of ionisation energy needed to remove an electron. Electron shielding, atomic radii and the nuclear charge of the atom.

Electron shielding is where electrons in the inner energy levels ‘shield’ the positive charge of the nucleus, so other electrons are not as strongly attracted to the nucleus. So, the larger the atomic radii= more electron shielding= less ionization energy needed to remove each electron.

Ionisation energies are also shown in equations, which you must be able to do!

Electronegativity Electronegativity is the ability of an atom in a covalent

bond to attract a bonded pair of electrons towards itself.

Nitrogen, Oxygen and Fluorine are the most electronegative elements in the periodic table.Element Approx. electronegativityNitrogen 3Oxygen 3.5Fluorine 4

Electronegativity is measured using the Pauling scale (which was invented by the chemist Linus Pauling in 1932)

Dipoles and Polar/non-polar molecules A permanent dipole is a small

charge difference across a bond that is formed from a difference in electronegativities of the bonded atoms.

A non-polar molecule a symmetrical molecule, without an overall dipole.                 

A polar molecule is - a non-symmetrical molecule, with an overall dipole.

An overall dipole is formed when one of the atoms in the bond, has a higher electronegativity (as it has the most electrons, as shown in the picture)

The atom with a high electronegativity is shown by a delta negative symbol.

Molecular Shapes

The shape of a molecule is determined by the repulsion between bonded electrons and non-bonded electrons (lone pairs).

Lone electron pairs repel more than bonded pairs of electrons and give rise to distorted shapes.

Each lone pair repels by 2.5 degreesBy deducing the number of bonded electron pairs and lone pairs of electrons, the shape of a molecule may be predicted. BF3 is trigonal planar; CH4 and NH4

+ are tetrahedral; SF6 is octahedral; H2O is non-linear (V-shaped/bent); CO2 is linear and ammonia, NH3 as pyramidal.

Intermolecular Forces

An intermolecular force exists between molecules and may include hydrogen bonding, dipole-dipole or van der Waals’ forces.

Don’t confuse intermolecular forces for intramolecular forces, intramolecular forces occur inside of molecules.

Electronegativity is the ability of an atom in a covalent bond to attract a bonded pair of electrons towards itself.

Hydrogen bonding arises in molecules in which a hydrogen atom is bonded to an N, O or F atom.

Water molecules, and other substances consisting of hydrogen bonding, have anomalous properties as a result.

Anomalous properties of water include ice floating on water, water having a high boiling point and the high surface tension that water has. These are all due to hydrogen bonds.

Bonding and Physical Properties

These are the molecular shapes of molecules ->

How many bonding pairs and example

Name Bond angle

2, e.g. CO2 LINEAR 1803, e.g. BF3 TRIGONAL

PLANAR120

4, e.g. CH4 TETRAHEDRAL 1095, e.g. PF5 TRIGONAL

BIPYRAMIDAL120 and 90

6, e.g. SF6 OCTRAHEDRAL 90

Metals consist of a close-packed arrangement of positive ions, through which delocalized electrons move.

Metals are very good electrical conductors as a result of having mobile electrons.

Giant structures have high melting and boiling points due to strong chemical bonds acting throughout the structure.

Giant ionic structures conduct electricity when molten, and when dissolved in water due to mobile ions, not electrons.

Periodicity

When the elements are arranged in order of their atomic number, there is a regular repetition of physical and chemical properties.

Elements in the same group have similar chemical and physical properties.

In the Periodic Table, ionization energies increase moving across a period from left to right, and decrease moving down a group.

Electron structures, atomic radii, melting points and boiling points all show periodicity.

Group 2 Elements – the Alkaline Earth Metals

The alkaline earth metals belong to the S block on the periodic table

Group 2’s first ionisation energies decreases down the group These elements all react with water to form a solution of the

hydroxide and hydrogen gas. These elements react with oxygen to form an oxide.

Reactivity increases on descending the group because the outer two electrons are further from the nucleus and are less shielded.

Metal hydroxides are weak alkalis and typically have a pH between 8 and 11

Group 2 carbonates decompose with greater difficulty as the group is descended, to form the metal oxide and carbon dioxide gas.

As you move down group 2 the elements hydroxides solubility and alkalinity increases.

Group 2 elements become more reactive DOWN the group Group 2 carbonates decompose at higher temperatures down

the group

The hydroxides of group 2 elements become more soluble in water, and the resulting solutions become more alkaline as you move down the group.

Group 2 elements have high melting and boiling points They are light metals with low densities They form colourless compounds and are oxidised in

reactions

Group 2 reactions with oxygen Group 2 elements react vigorously with oxygen, this is a

redox reaction. The product is an ionic oxide with the general formula being

MO where the M stands for the group 2 element The group 2 element is oxidised in this process and looses 2

electrons Oxygen is reduced in this process and gains the 2 electrons

Group 2 reactions with water Group 2 elements react with water to form hydroxides

with the general formula M(OH)2. Hydrogen gas is also produced.

During the reaction of a group 2 element with water, only 1 of the Hydrogen atoms is reduced, the other hydrogen atom’s oxidation number remains the same, as +1.

As you move down group 2 the metals react more vigorously with

Uses of group 2 hydroxides Calcium hydroxide is used by farmers and gardeners as ‘lime’

to neutralise acidic soil. Magnesium hydroxide is used as milk of magnesia to relieve

indigestion. It works by neutralising any excess stomach acid.

Group 7 Elements – the Halogens

All halogens exist as diatomic molecules in which van der Waals’ intermolecular forces act between the molecules.

Halogens dissolve in organic solvents, like hexane, to form characteristic colours, for example, iodine forms a purple solution.

Halogen atoms gain one electron to form halide ions, X-, and this ability becomes easier on moving up the group.

Halogen atoms become larger on descending the group, so a gained electron is only weakly attracted due to greater shielding.

Halides act as oxidising agents, as they accept/gain an electron from another element, so the halide is reduced and the other element is oxidised.

Halides belong to the P block on the periodic table as they fill the P sub-shell.

The boiling points of the halogens increase down the group as the number of electrons increases so there are more van der Waal’s forces to break, which takes energy in the form of heat.

Fluorine is the most reactive element in group 7 as it has fewer shells, electrons, electron shielding and smaller nuclear radii so there is more attraction to the outer electrons, making it easier to form an ion.

Astatine is the least reactive element in group7 as it has more shells, electrons, electron shielding and larger nuclear radii so there is less attraction to the outer electrons making it harder to form ions.

More reactive halides will displace less reactive halides in reactions.

Redox reactions of the halogens Redox reactions of the halogens show that halogens become

less able to form halide ions as you move down the group, this is done by reacting aqueous solutions of halide ions Cl- (aq), Br- (aq) and I- (aq) with aqueous solutions of the halogens Cl2 (aq) , Br2 (aq) and I2 (aq)

Each halogen is mixed with an aqueous solution of the different halides. A more reactive halogen will oxidise and displace a halide of a less reactive halogen. This is a displacement reaction

As halogens form solutions with distinct colours it is easy to see if a displacement reaction has taken place.

Only more reactive halogens displace less reactive halogens.

As chlorine is higher up group 7, so its more reactive, it will displace and oxidise both Br- and I- ions

Bromine only oxidises and displaces Iodine, and iodine does not oxidise either Cl- or Br- ions.

Reactions of Chlorine and Halide Ions

Disproportionation is where the same element is oxidised and reduced in a reaction.

For example, this happens when chlorine is added to waterCl2 + H2o -> HCl + HClO. Here Cl is oxidised and reduced.

Chlorine disproportionates in water to form hydrochloric acid and chloric(I) acid, the latter being an oxidising agent.

Chlorine also disproportionates in cold, aqueous sodium hydroxide to form sodium chloride, sodium chlorate(I) and water.

Cl2(aq) +2NaOH(aq) ->NaCl(aq) + NaClO(aq) +H2O (l)

Chlorine is used to kill bacteria in water supplies, but is also toxic to humans at higher doses.

Halide ions are detected with silver(I) nitrate solution and the subsequent reaction with ammonia solution.

Halide ions have distinct precipitate colours when mixed with cyclohexane.

Halogen Colour in water

Colour in cyclohexane

Cl2 Pale green Pale-greenBr2 orange orangeI2 brown violet

Questions

Exercise 1 – Avogadro’s Number and reacting masses

Avogadro’s Number

1. Calculate the number of particles in the following substances:a) 0.025 molesb) 2.5 g of CO2c) 5.0 g of Pbd) 100 g of N2

2. Calculate the mass of the following:a) 2.5 x 1023 molecules of N2b) 1.5 x 1024 molecules of CO2c) 2 x 1020 atoms of Mg

Reacting Masses

3. Calculate the mass of H2O required to react completely with 5.0 g of SiCl4:

SiCl4 + 2H2O SiO2 + 4HCl

4. Calculate the mass of phosphorus required to make 200 g of phosphine, PH3, by the reaction:P4(s) + 3NaOH(aq) +3H2O(l) 3NaH2PO4(aq) + PH3(g)

5. Lead (IV) oxide reacts with concentrated hydrochloric acid as follows:PbO2(s) + 4HCl(aq) PbCl2(s) + Cl2(g) + 2H2O(l) What mass of lead chloride would be obtained from 37.2g of PbO2, and what mass of chlorine gas would be produced?

6. When copper (II) nitrate is heated, it decomposes according to the following equation:2Cu(NO3)2(s) 2CuO(s) + 4NO2(g) + O2(g).

When 20.0g of copper (II) nitrate is heated, what mass of copper (II) oxide would be produced? What mass of NO2 would be produced?

Exercise 2 -Using molarities and concentrations

1. Calculate the number of moles of H2SO4 in 50 cm3 of a 0.50 moldm-3 solution.

2. Calculate the number of moles of FeSO4 in 25 cm3 of a 0.2 moldm-3 solution.

3. Calculate the mass of KMnO4 in 25 cm3 of a 0.02 moldm-3 solution.

4. Calculate the mass of Pb(NO3)2 in 30 cm3 of a 0.1 moldm-3 solution.

5. What is the molarity of 1.06g of H2SO4 in 250 cm3 of solution?

6. What is the molarity of 15.0 g of CuSO4.5H2O in 250 cm3 of solution?

7. How many moles of NaCl are there in 25 cm3 of a 50 gdm-3 solution?

Exercise 3-Reacting masses and solutions

1. 25 cm3 of a solution of 0.1 moldm-3 NaOH reacts with 50 cm3 of a solution of hydrochloric acid. What is the molarity of the acid?

2. 25.0 cm3 of a 0.10 moldm-3 solution of sodium hydroxide was titrated against a solution of hydrochloric acid of unknown concentration. 27.3 cm3 of the acid was required. What was the concentration of the acid?

3. A solution of hydrochloric acid of volume 25.0 cm3 was pipetted onto a piece of marble which is calcium carbonate. When all action

had ceased, 1.30g of the marble had dissolved. Find the concentration of the acid Equation: CaCO3 + 2HCl CaCl2 + CO2 + H2O

4. What volume of 0.1 moldm-3 hydrochloric acid would be required to dissolve 2.3 g of calcium carbonate?Equation: CaCO3(s) + 2HCl(aq) CaCl2(aq) + CO2(g) + H2O(l)

5. 2.05 g of the carbonate of an unknown alkali metal (X2CO3) required 8.9 cm3 of 2.0 moldm-3 hydrochloric acid to completely dissolve it. What was the relative atomic mass of the metal and which metal was it?Equation: X2CO3(s) + 2HCl(aq) 2XCl(aq) + CO2(g) + H2O(l)

Exercise 4 - empirical and molecular formulae

1. A compound contains C 62.08%, H 10.34% and O 27.58% by mass. Find its empirical formula and its molecular formula given that its relative molecular mass is 58.

2. Find the empirical formula of the compound containing C 22.02%, H 4.59% and Br 73.39% by mass.

3. A compound containing 85.71% C and 14.29% H has a relative molecular mass of 56. Find its molecular formula.

4. A compound containing 84.21% carbon and 15.79% hydrogen by mass has a relative molecular mass of 114. Find its molecular formula.

5. Analysis of a hydrocarbon showed that 7.8 g of the hydrocarbon contained 0.6 g of hydrogen and that the relative molecular mass was 78. Find the molecular formula of the hydrocarbon.

6. 3.36 g of iron join with 1.44 g of oxygen in an oxide of iron. What is the empirical formula of the oxide?

7. What is the percentage composition of SiCl4?

8. What is the mass of sulphur in 1 tonne of H2SO4?

Exercise 5 – Water of Crystallisation

1. A sample of hydrated calcium sulphate, CaSO4.xH2O, has a relative formula mass of 172. What is the value of x?

2. A hydrated salt is found to have the empirical formula CaN2H8O10. What is its dot formula?

3. A hydrated carbonate of an unknown Group 1 metal has the formula X2CO3.10H2O and is found to have a relative formula mass of 286. What is the Group 1 metal?

4. 11.25 g of hydrated copper sulphate, CuSO4.xH2O, is heated until it loses all of its water. Its new mass is found to be 7.19 g. What is the value of x?

5. 13.2 g of a sample of zinc sulphate, ZnSO4.xH2O, was strongly heated until no further change in mass was recorded. On heating, all the water of crystallisation evaporated as follows: ZnSO4.xH2O ZnSO4 + xH2O.Calculate the number of moles of water of crystallisation in the zinc sulphate sample given that 7.4 g of solid remained after strong heating.

6. A sample of hydrated magnesium sulphate, MgSO4.xH2O, is found to contain 51.1% water. What is the value of x.

Exercise 6 – titration calculations

1. 25 cm3 of a solution of 0.1 moldm-3 NaOH reacts with 50 cm3 of a solution of hydrochloric acid. What is the molarity of the acid?

2. 25.0 cm3 of a 0.10 moldm-3 solution of sodium hydroxide was titrated against a solution of hydrochloric acid of unknown concentration. 27.3 cm3 of the acid was required. What was the concentration of the acid?

3. 25 cm3 of a solution of sodium hydroxide reacts with 15 cm3 of 0.1 mol/dm3 HCl. What is the molar concentration of the sodium hydroxide solution?

4. Succinic acid has the formula (CH2)n(COOH)2 and reacts with dilute sodium hydroxide as follows: (CH2)n(COOH)2 + 2NaOH (CH2)n(COONa)2

+ 2H2O2.0 g of succinic acid were dissolved in water and the solution made up to 250 cm3. This solution was placed in a burette and 18.4 cm3 was required to neutralise 25 cm3 of 0.1 moldm-3 NaOH. Deduce the molecular formula of the acid and hence the value of n.

5. Sodium carbonate exists in hydrated form, Na2CO3.xH2O, in the solid state. 3.5 g of a sodium carbonate sample was dissolved in water and the volume made up to 250 cm3. 25.0 cm3 of this solution was titrated against 0.1 moldm-3 HCl and 24.5 cm3 of the acid were required. Calculate the value of x given the equation:Na2CO3 + 2HCl 2NaCl + CO2 + H2O

6. 25 cm3 of a sample of vinegar (CH3COOH) was pipetted into a volumetric flask and the volume was made up to 250 cm3. This solution was placed in a burette and 13.9 cm3 were required to neutralise 25 cm3 of 0.1 moldm-3

NaOH. Calculate the molarity of the original vinegar solution and its concentration in gdm-3, given that it reacts with NaOH in a 1:1 ratio.

7. 2.5 g of a sample of ethanedioic acid, H2C2O4.nH2O, was dissolved in water and the solution made up to 250 cm3. This solution was placed in a burette and 15.8 cm3 were required to neutralise 25 cm3 of 0.1 moldm-3 NaOH. Given that ethanedioic acid reacts with NaOH in a 1:2 ratio, calculate the value of n.

These are all the definitions you need to know for the exam. Print them off, cut along each row, then stick them back to back to use as a revision activity!

1st ionisation energyThe amount of energy needed to remove 1 electron, from each atom, in a mole of atoms in the gas phase.

Successive ionisation energiesA measure of the amount of energy required to remove each electron in turn.

IsotopeAn atom of an element with the same number of protons but a different number of neutrons

Relative atomic massThe weighted mean mass of an atom of an element compared to the weight of 1/12th of an atom of carbon-12

Relative isotopic massThe mass of an atom of an isotope compared to 1/12th of the mass of an atom of carbon-12

OrbitalA region of space around a nuclei of an atom that can hold 2 electrons in opposite spins

Electron configuration The arrangement of electrons in an atom

Relative formula massThe weighted mean mass of a formula unit compared to 1/12th of the mass of carbon-12

Molar massThe mass per mole of a substance. Which is measured in g/mol

SubshellA group of the same type of atomic orbitals inside of the same electron shell

MoleA mole of a substance contains the same amount of particle that there are in 12.0g of carbon-12 (6.023 x 1023)

Empirical formulaThe simplest whole number ratio of atoms of an element in a compound

MoleculeA small group of molecules held together by covalent bonds

Molecular formulaThe actual number of atoms of an element in a compound

Ionic bondingThe electrostatic attraction between oppositely charged ions

Giant ionic latticeElectrostatic forces of attraction between oppositely charged ion, which extend into 3D

Simple molecular latticeA 3D structure of molecules, that are bonded together by weak intermolecular forces

Covalent bondA bond that is formed from a shared pair of electrons

Dative covalent bonding/ co-ordinate bonding

A shared pair of electrons that came from the same atom

permanent dipolea small charge difference across a bond that is formed from a difference in electronegativities of the bonded atoms

Permanent dipole-dipole forceA weak attractive force between permanent dipole in neighbouring polar molecules

Polar covalent bondA covalent bond that has a permanent dipole

Polar molecule A molecule with an overall dipole

Electronegativity the ability of an atom in a covalent bond to attract a bonded pair of electrons towards itself

Metallic bondA bond formed when positive metal cations are attracted to negatively charged delocalised electrons.

Hydrogen bondingthe electromagnetic attractive interaction of a hydrogen atom and an electronegative atom, such as nitrogen, oxygen or fluorine.

Van der Waals forcesWeak electrostatic attractive forces between uncharged molecules, arising from the interaction of permanent dipoles

Intermolecular forceAn attractive force between neighbouring molecules.

Acid A proton donor in a solution

Base A proton acceptor in a solution

AlkaliA base that is dissolved in water, forming hydroxide ions,OH-

OxidationThe loss of electrons or an oxidation number increase

ReductionThe gain of electrons or an oxidation number decrease

Oxidising agent cause oxidation by accepting an electron

Reducing agent Cause reduction by donating an electron

Redox reactionA reaction where both oxidation and reduction take place

SaltAn ionic compound that results from the neutralization reaction of an acid and a base.

PeriodicityThe repetition of similar properties in chemical elements, as indicated by their positioning in the periodic table.

PeriodA horizontal row of elements in the periodic table

GroupA vertical column of element in the periodic table

Thermal decomposition Chemical decomposition caused by heat.

Aqueous A solution in which the solvent is water

Molar volume The volume per mole of a gas.

ConcentrationThe amount of solute in moles dissolved per 1 dm3

Standard solution A solution with a known concentration

stoichiometryThe molar relationship between the relative quantities of substances taking part in a reaction.

Cation Positively charged ion

Anion Negatively charged ion

Hydrated saltA crystalline substance containing water molecules.

Anhydrous saltA crystalline substance that does not contain water molecules.

Water of crystallisationWater molecules that form an essential part of the crystalline structure of a compound.

TitrationA method of determining a concentration of a solution.

DisproportionationWhere an element is oxidised and reduced in the same reaction.

Displacement reactionA reaction where a more reactive element displaces a less reactive element, from an aqueous solution of the latter ion’s

OxyanionsNegative ions that contain an element along with oxygen.

SpeciesAny type of particle that takes place in a chemical reaction

Oxidation numberIs a measure of the number of electrons that an atom uses to bond with atoms of another element

Electron shieldingThe repulsion between electrons in different inner shells.

CompoundA substance formed from 2 or more chemically bonded elements in a fixed ratio.

Lone pair of electronsAn outer-shell pair of electrons that are not involved in chemical bonding.

Precipitation reactionThe forming of a solid from an aqueous solution during a chemical reaction.