key

Question 1

John Dalton (1803):
Atom model

Answer 1

● He published his own three-part atomic theory:
1. All substances are made of atoms. Atoms are small particles that cannot be created, divided, or destroyed.

2. Atoms of the same element are exactly alike, and atoms of different elements are different.
3. Atoms join with other atoms to make new substances.

● Much of Dalton’s theory was correct, but some of it was later proven incorrect and revised as scientists learned more about atoms

Question 2

J.J. Thomson (1897):

Plum pudding model

Answer 2

● Used a cathode-ray tube to conduct an experiment
● This discovery identified an error in Dalton’s atomic theory. Atoms can be divided into smaller parts.
● Because the beam moved away from the negatively charged plate and toward the positively charged plate, Thomson knew that the particles must have a negative charge.
● Thomson proposed a model of an atom called the “plum-pudding” model, in which negative electrons are scattered throughout soft blobs of positively charged material

Question 3

Ernest Rutherford (1909):
Atom

Answer 3

● Shot a beam of positively charged particles into a sheet of gold foil.
● Most of the particles did continue in a straight line (as you would expect from plum pudding model). However some of the particles were deflected to the sides
a bit, and a few bounced straight back.
● Rutherford developed a new model which said that most of the atom’s mass is found in a region in the center called the nucleus.
● In Rutherford’s model the atom is mostly empty space, and the electrons travel in random paths around the nucleus

Question 4

Structure of an atom

Answer 4

A nucleus containing protons and neutrons, surrounded by electrons in shells

Question 5

Relative Charge and Relative mass:

Proton

Answer 5

Relative charge = +1

Relative mass = 1

Question 6

Relative Charge and Relative mass:

Neutron

Answer 6

Relative charge = 0

Relative mass = 1

Question 7

Relative Charge and Relative mass:

Electron

Answer 7

Relative charge = -1

Relative mass = 1/1836

Question 8

Why do atoms contain equal number of protons and electrons

Answer 8

● Atoms are neutral and the charges on a proton are +1 and on an electron are -1
● therefore amount of protons = amount of electrons, so that the charges cancel

Question 9

Describe the nucleus of an atom compared to the

overall size of the atom

Answer 9

The nucleus of an atom is very small compared to the overall size of the atom

Question 10

Where is most of the mass in an atom concentrated?

Answer 10

Most of the mass of an atom is concentrated in the

nucleus

Question 11

What is the meaning of the term mass number of an atom?

Answer 11

● Mass (nucleon) Number = number of protons + neutrons

Question 12

How would you describe atoms of a given element?

Answer 12

As having the same number of protons in the nucleus and that this number is unique to that element

Question 13

Isotopes

Answer 13

Isotopes are different atoms of the same element containing the same number of protons but different numbers of neutrons in their nuclei

Question 14

How would you calculate the numbers of protons, neutrons and electrons in atoms given the atomic number and mass number?

Answer 14

Atomic (proton) Number = number of protons (= number of electrons if it’s an atom, because atoms are neutral)
● therefore, you can calculate number of neutrons by doing mass number - atomic number

Question 15

Explain how the existence of isotopes results in relative atomic masses of some elements not being whole numbers

Answer 15

● because isotopes have the same number of protons but different numbers of neutrons, they are still atoms of the same element, but they have different atomic masses
● the relative atomic mass is calculated using the abundance of different isotopes and because it is an average it can lead to the relative atomic mass not being a whole number (atomic number and mass number will always be whole numbers - they are not averages)

Question 16

How do we describe mass of an atom?

Answer 16

● since the mass of atoms is so small, we compare their masses to each other. A carbon atom having a mass number 12, i.e. (12C) is taken as standard for this
comparison and its relative atomic mass is 12.
● It is written as A​r​ or R.A.M.​.

Question 17

How to calculate the relative formula mass:

Answer 17

1. Write the formula of the compound.
2. Write the numbers of each atom in the formula.
3. Insert the relative atomic mass for each type of atom.
4. Calculate the total mass for each element.
5. Add up the total mass for the compound.

Question 18

A sample of chlorine gas is a mixture of 2 isotopes,
chlorine-35 and chlorine-37. These isotopes occur in
specific proportions in the sample i.e. 75% chlorine-35
and 25% chlorine-37. Calculate the R.A.M. of chlorine in
the sample.
Calculate the RAM of chlorine in this sample

Answer 18

R.A.M. =
​(mass of isotope-A x % of isotope-A) + (mass of isotope-B x % of isotope-B) /​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​​100
​​ =
(35 x 75) + (37 x 25)/​​​​​​​​100
=
3550/100
R.A.M. = 35.5

Question 19

Describe how Mendeleev arranged the elements, known at that time, in a periodic table by using properties of these elements and their compounds

Answer 19

● He ordered his table in order of atomic mass, but not always strictly – i.e. in some places he changed the order based on atomic weights.
● Left gaps for elements that he thought had not been discovered yet.

Question 20

Describe how Mendeleev used his table to predict the existence and properties of some elements not then discovered

Answer 20

● Mendeleev realised elements with similar properties belonged in the same groups in the periodic table so was able to leave gaps and place the discovered
elements where they fit best
● Elements with properties predicted by Mendeleev were later discovered and filled the gaps

Question 21

Explain that Mendeleev thought he had arranged elements in order of increasing relative atomic mass but this was not always true

Answer 21

● Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct, because some elements have a higher mass than others when isotopes are taken into account, but a lower one if you only look at one specific isotope.

Question 22

Explain the meaning of atomic number of an element in terms of position in the periodic table and number of protons in the nucleus

Answer 22

● Elements are arranged in order of atomic (proton) number (bottom number) and so that elements with similar properties are in columns, known as groups.
● Elements in the same group have the same amount of electrons in their outer shell, which gives them similar chemical properties.

Question 23

How are elements arranged in the periodic table?

Answer 23

●Elements are arranged in order of increasing atomic number, in rows called periods and elements
●Elements with similar properties are placed in the same vertical columns called groups

Question 24

Metals in periodic table

Answer 24

● Metals = elements that react to form positive ions.
o Majority of elements are metals.
o Found to the left and towards the bottom of the periodic table., because they lose electron(s) in order to form these positive ions, forming an electronic structure that is stable, like that of a noble gas

Question 25

Non metals in periodic table

Answer 25

● Non-metals = elements that do not form positive ions. o Found towards the right and top of the periodic table, because they gain electron(s) in order to form these negative ions, forming an electronic structure that is stable, like that of a noble gas

Question 26

How to represent electronic configurations of an atom?

Answer 26

● the electronic configuration of an element tells you how many electrons are in each shell around an electron’s nucleus ● for example, sodium has 11 electrons: 2 in its most inner shell, then 8, then 1 in its outermost shell. ○ you can represent sodium’s electronic configuration as: 2.8.1

Question 27

Explain how the electronic configuration of an element is related to its position in the periodic table

Answer 27

● the group an electron is in tells you how many electrons are in its outermost shell aka group 1 elements have 1 electron in their outer shell ● the period an electron is in tells you which number shell an element's outermost electron is found in aka period 3 elements have their outermost electrons in shell 3 ● remember all the shells up until the shell will be full (for the 1st shell this means 2 electrons and for shells 2 and 3 this means 8 electrons)

Question 28

Explain how ionic bonds are formed

Answer 28

● Metals + nonmetals: electrons in the outer shell of the metal atom are transferred o Metal atoms lose electrons to become positively charged ions (cation) o Nonmetal atoms gain electrons to become negatively charged ions (anion)

Question 29

What is an ion

Answer 29

● Since an ion is formed from a metal losing an electron, i.e. becoming a positive metal ion or from a non-metal gaining an electron, i.e. becoming a negative ion…​ An ion is an atom or group of atoms with a positive or negative charge

Question 30

How would you calculate the numbers of protons, neutrons and electrons in simple ions given the atomic number and mass number?

Answer 30

● Atomic number = proton number = number of protons ● Mass number = nucleon number = number of protons + neutrons ● In an atom number of protons = number of electrons, but in an ion, there is a different number of electrons to protons. to work out electrons in an ion: o work out how many electrons an atom of the element would have (same as proton number) o work out how many electrons have been lost or gained (using charge remember -ve means electrons gained, +ve means electrons lost) o calculate number of electrons in atom plus electrons gained or minus electrons lost

Question 31

Explain the formation of ions in ionic compounds from their atoms

Answer 31

● Ions produced by metals in Groups 1 and 2 and by nonmetals in Groups 6 and 7 have the electronic structure of a noble gas (Group 0) ● this means group 1 metals will lose 1 electron and form +1 ions ● group 2 metals will lose 2 electrons and form +2 ions ● group 6 nonmetals will gain 2 electrons and form 2- ions ● group 7 nonmetals will gain 1 electron and form 1- ions ● remember a compound will have an overall charge of 0 so you need to balance out the + and - charges

Question 32

Explain the use of the endings –ide and –ate in the names of | compounds

Answer 32

● these endings are used for the negatively charged ions in a compound ● -ide means the compound contains 2 elements (one is the nonmetal -ve ion) ● -ate means the compound contains at least 3 elements, one of which is oxygen

Question 33

How would you deduce the formula of ionic compounds?

Answer 33

To deduce the formula of ionic compounds, you need to balance out the + and - charges to make the overall charge 0. You do this by writing a little number below the element e.g. Cl3 or for ions with more than one element you draw a bracket round first e.g. (SO4)2

Question 34

Explain the structure of an ionic compound

Answer 34

● A giant structure of ions = ionic compound ● Held together by strong electrostatic forces of attraction between oppositely charged ions ● The forces act in all directions in the lattice, and this is called ionic bonding. ● The lattice has a regular arrangement of ions

Question 35

How is a covalent bond formed?

Answer 35

● Covalent bonding occurs in most non-metallic elements and in compounds of nonmetals ● When atoms share pairs of electrons, they form covalent bonds. These bonds between atoms are strong.

Question 36

Recall that covalent bonding results in the formation of molecules

Answer 36

● Covalently bonded substances may consist of small molecules e.g. HCl,H2,O2,Cl2, NH3, CH4 .● Some have very large molecules, such as polymers. ● Some have giant covalent structures (macromolecules) e.g diamond, silicon dioxide.

Question 37

Recall the typical size (order of magnitude) of atoms and small molecules

Answer 37

● Simple molecular substances consist of molecules in which the atoms are joined by strong covalent bonds ● Therefore, atoms are smaller than small molecules

Question 38

Properties of ionic compounds

Answer 38

● Ionic compounds are made up of a metal and a nonmetal ● Ionic compounds have regular structures (giant ionic lattices) in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions. ● They have high melting and boiling points, because a lot of energy is required to break the many strong bonds. ● When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and carry current, and they do not conduct electricity as solids, because the ions are fixed and are not able to move, carrying charge with them. ● Often dissolve in water to form an aqueous solution

Question 39

Properties of simple molecular compounds

Answer 39

● Substances that consist of small molecules are usually gases or liquids that have low boiling and melting points. They are made up of nonmetal elements. ● Substances that consist of small molecules have weak intermolecular forces between the molecules. These are broken in boiling or melting, not the covalent bonds. ○ The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points. ● Substances that consist of small molecules don’t conduct electricity, because small molecules do not have an overall electric charge. although, some breakdown in water to form ions which can conduct electricity ● Many are insoluble in water, but some are soluble because they can form intermolecular forces with water which are stronger than those between water molecules or their own molecules already (e.g. CO2 and NH3 are soluble)

Question 40

Giant Covalent Structures

Answer 40

● They are made up of nonmetal elements ● Substances that consist of giant covalent structures are solids with very high melting points. o All of the atoms in these structures are linked to other atoms by strong covalent bonds. ▪ These bonds must be overcome to melt or boil these substances. ● some giant covalent structures can conduct electricity, whereas others can’t

Question 41

Properties of metals

Answer 41

● Metals consist of giant structures of atoms arranged in a regular pattern. They are always made up of just metallic elements ● The electrons in the outer shell of metal atoms are delocalised and so are free to move through the whole structure. ● The sharing of delocalised electrons gives rise to strong metallic bonds.

Question 42

Structure of metals

Answer 42

● Metals have giant structures of atoms with strong metallic bonding. o Therefore, most metals have high melting and boiling points. o They can conduct heat and electricity because of the delocalised electrons in their structures. o Conduction depends on the ability for electrons to move throughout the metal. o The layers of atoms in metals are able to slide over each other, so metals can be bent and shaped. o insoluble in water- but some will react with it instead

Question 43

properties of ionic compounds

Answer 43

● high melting and boiling points- strong electrostatic forces ● conduct electricity when molten/dissolved- ions can move ● don’t conduct electricity when solid- ions are fixed in place

Question 44

properties of typical covalent, simple molecular compounds

Answer 44

● low melting and boiling points- only weak forces between molecules must be overcome, not covalent bonds ● poor conduction of electricity- no charged particles or electrons that are free to move

Question 45

What are graphite and diamond examples of?

Answer 45

Graphite and diamond are different forms of carbon and they are examples of giant covalent substances

Question 46

Structure of diamonds

Answer 46

● In diamond, each carbon is joined to 4 other carbons covalently. o It’s very hard, has a very high melting point and does not conduct electricity.

Question 47

Structure of graphite

Answer 47

● In graphite, each carbon is covalently bonded to 3 other carbons, forming layers of hexagonal rings, which have no covalent bonds between the layers. o The layers can slide over each other due to no covalent bonds between the layers, but weak intermolecular forces. Meaning that graphite is soft and slippery. ● One electron from each carbon atom is delocalised. o This makes graphite similar to metals, because of its delocalised electrons. o It can conduct electricity – unlike diamond.

Question 48

Diamond uses

Answer 48

o Cutting tools – very hard, due to its rigid structure

Question 49

Graphite uses

Answer 49

o Electrodes – graphite can conduct electricity – unlike Diamond o Lubricant – weak intermolecular forces and no covalent bonds between the layers, therefore it is soft and slippery

Question 50

Properties of fullerenes: | Graphene

Answer 50

o Single layer of graphite | o Has properties that make it useful in electronics and composites

Question 51

Fullerene

Answer 51

Molecules made up entirely of carbon, in the shape of a sphere or a tube

Question 52

Properties of fullerenes: | Carbon

Answer 52

Carbon can also form fullerenes with different numbers of carbon atoms. o Molecules of carbon atoms with hollow shapes o They are based on hexagonal rings of carbon atoms, but they may also contain rings with five or seven carbon atoms o The first fullerene to be discovered was Buckminsterfullerene (C60), which has a spherical shape

Question 53

Properties of fullerenes: | Carbon nanotubes

Answer 53

o Cylindrical fullerenes with very high length to diameter ratios o Their properties make them useful for nanotechnology, electronics and materials

Question 54

Examples of uses

Answer 54

o They can be used as lubricants, to deliver drugs in the body and catalysts. o Nanotubes can be used for reinforcing materials, for example tennis rackets.

Question 55

Describe how simple polymers | consist of large molecules containing chains of carbon atoms

Answer 55

● Polymers have very large molecules (the n in the diagram structure shows there are many many repeat units) ● Atoms in the polymer molecules are linked to other atoms by strong covalent bonds ● Intermolecular forces between polymer molecules are relatively strong and so these substances are solids at room temperature

Question 56

properties of metals, including malleability and the ability | to conduct electricity

Answer 56

● malleable- the layers of atoms in metals are able to slide over each other ● can conduct electricity- delocalised electrons can move

Question 57

Describe the limitations of particular representations and models

Answer 57

● Main limitation is that it applies really well only to the small class of solids composed of Group 1 and 2 elements with highly electronegative elements such as the halogens ● In covalent molecular, the dot-cross diagrams don’t express the relative attraction of shared electrons due to electronegativity ● 2d diagrams don’t show the 3d arrangement of atoms, and 3d diagrams don’t show the share or transfer of electrons

Question 58

Metals

Answer 58

Metals as shiny solids which have high melting points, high density and are good conductors of electricity whereas most non-metals have low boiling points and are poor conductors of electricity

Question 59

How to calculate relative formula mass given relative atomic masses

Answer 59

● Relative formula mass (Mr) of a compound: sum of the relative atomic masses of the atoms in the numbers shown in the formula (remember you could have more than 1 atom of a certain element in a compound e.g. in CaCl2, there are 2 atoms of chlorine so you need to add on 35.5 x2) ● In a balanced chemical equation: sum of Mr of reactants in quantities shown = sum of Mr of products in quantities shown

Question 60

Calculate the formulae of simple compounds from reacting masses

Answer 60

a. work out moles of each using moles = mass ¨ molar mass b. work out the ratio of moles c. times the ratio so that you get the smallest whole numbers possible d. find the formula by timesing each element by their number in the ratio (remember to use little numbers not a big number at the front) ● This is an empirical formula because it shows the simplest ratio of the number of atoms of different types of elements in a compound

Question 61

● Empirical formula from the formula of molecule:

Answer 61

o if you have a common multiple e.g. Fe2O4, the empirical formula is the simplest whole number ratio, which would be FeO2 o if there is no common multiple, you already have the empirical formula

Question 62

Molecular formula from empirical formula and relative molecular mass

Answer 62

o Find relative molecular mass of the empirical formula o Divide relative molecular mass of compound by that of the empirical formula o Multiply the number of each type of atom in the empirical formula by this number o e.g. if answer was 2 and the empirical formula was Fe2O3 then the molecular formula would be empirical formula x 2 = Fe4O6

Question 63

Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide METHOD

Answer 63

● weigh some pure magnesium ● Heat magnesium to burning in a crucible to form magnesium oxide, as the magnesium will react with the oxygen in the air ● weigh the mass of the magnesium oxide ● Known quantities: mass of magnesium used & mass of magnesium oxide produced

Question 64

Describe an experiment to determine the empirical formula of a simple compound such as magnesium oxide ● Required calculations

Answer 64

○ mass oxygen = mass magnesium oxide - mass magnesium ○ moles magnesium = mass magnesium ÷ molar mass magnesium ○ moles oxygen = mass oxygen ÷ molar mass oxygen ○ calculate ratio of moles of magnesium to moles of oxygen ○ use ratio to form empirical

Question 65

Law of conservation of mass

Answer 65

● Law of conservation of mass: no atoms are lost or made during a chemical reaction so the mass of the products = mass of the reactants o Therefore, chemical reactions can be represented by symbol equations, which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation.

Question 66

Law of conservation of mass applied to a precipitation reaction in a closed flask

Answer 66

With a precipitation reaction – precipitate that forms is insoluble and is a solid, as all the reactants and products remain in the sealed reaction container then it is easy to show that the total mass is unchanged

Question 67

Law of conservation of mass applied to a reaction in an open flask

Answer 67

Does not hold for a reaction in an open flask that takes in or gives out a gas, since mass will change from what it was at the start of the reaction as some mass is lost when the gas is given off

Question 68

carbon dioxide

Answer 68

co2

Question 69

water

Answer 69

H20

Question 70

Oxygen

Answer 70

O2

Question 71

Hydrogen

Answer 71

H2

Question 72

Nitrogen

Answer 72

N2

Question 73

Ammonia

Answer 73

NH3

Question 74

Hydrochloric acid

Answer 74

HCL

Question 75

Sulfuric acid

Answer 75

H2SO4

Question 76

Formula for Carbonate ion

Answer 76

CO3^2-