topic 1 - key concepts in chemistry

describe how the dalton model of an atom has changed over time

• the dalton model of an atom was a solid sphere containing no sub-atomic particles

• as sub-atomic particles were discovered, they were added to the dalton model to become the ‘plum pudding model'

  • plum pudding model was a solid, positively charged sphere with negatively charged electrons interspersed throughout

• rutherford model - over time, electrons were added to shells surrounding the proton and neutron nucleus

describe the structure of an atom

• a positively-charged nucleus

• containing protons and neutrons

• surrounded by negatively-charged electrons

• in shells

recall the relative charge of a proton

+1

recall the relative mass of a proton

1

recall the relative charge of a neutron

0

recall the relative mass of a neutron

1

recall the relative charge of an electron

-1

recall the relative mass of an electron

0.0005

explain why atoms contain equal numbers of electrons and protons

• atoms have a neutral charge

• an equal number of protons and electrons cancel out any net positive or negative charge

• causing the atom to remain neutrally-charged

describe the size of the nucleus of an atom compared to the overall size of the atom

very small in comparison

recall where the majority of mass is concentrated in an atom

nucleus

state the meaning of mass number of an atom

the sum of the number of protons and neutrons in the nucleus

describe what different atoms of the same given element have

• same number of protons in the nucleus

• the number of protons in the nucleus is unique to the given element

describe what isotopes are

• different atoms of the same element

• containing the same number of protons

• but a different number of neutrons in the nucleus

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

• isotopes of the same element have the same number of protons but different numbers of neutrons to each other

• this leads to different atomic masses with each isotope

• when finding the mean average of these results, the relative atomic mass may appear as a decimal as this is the average of the isotope’s atomic masses

state how to calculate relative atomic mass of an element from given isotopes

(mass of i1 x abundance of i1) + (mass of i2 x abundance of i2) / 100

describe how mendeleev arranged the known elements at the time into a periodic table

• columns - similar chemical properties of elements and their compounds

• rows - increasing atomic mass

• gaps - left gaps for undiscovered elements

• exceptions - there were a few exceptions, like iodine, which didn’t fit the pattern by atomic mass

state why mendeleev did not use relative atomic mass in his periodic table

because isotopes had not yet been discovered, meaning relative atomic masses of elements could not be calculated

describe how mendeleev used his periodic table to predict the existence and properties of undiscovered elements

• mendeleev used gaps in his table as place-holders for undiscovered elements

• due to the location of these gaps within the table, mendeleev could calculate their general chemical properties and atomic masses

• due to trends found within the columns and rows of the table

explain the issues with mendeleev’s table due to undiscovered isotopes

• mendeleev thought he had organised the elements into correct ascending atomic mass within their rows

• but calculating relative atomic mass was not taken into consideration as isotopes had not been discovered

• leading to the final atomic mass being wrong

• and the element ending up in the wrong space in the table

explain how elements are arranged by atomic number in the modern periodic table

• the modern periodic table is organised in increasing consecutive atomic number

• in periods

describe the organisation of the modern periodic table

• rows (periods) - elements are arranged in increasing atomic number

• columns (groups) - elements are arranged in groups of similar properties

state the definition of atomic number

number of protons in the nucleus of an atom

state the definition of atomic mass

number of protons and neutrons in the nucleus of an atom

state how to determine if an element is a metal based on its location in the periodic table

metal elements are located from the left to the end of the middle in the periodic table

state how to determine if an element is a non-metal based on its location in the periodic table

non-metal elements are located in the top right of the periodic table

explain why there is a division between metals and non-metals in the periodic table due to atomic structure (metals)

• the metal elements further to the left of the table have less electrons in their outer shells

• as you descend the groups of metals, the outer shell electrons become further away from the nucleus due to increasing atomic size

• this causes an increase in reactivity as you descend the groups

• making it more likely for the elements to lose electrons during a reaction

explain why there is a division between metals and non-metals in the periodic table due to atomic structure (non-metals)

• the non-metal elements further to the right have more electrons in their outer shells

• as you descend the groups of non-metals, there is a decrease in reactivity

• making it more likely for the elements to gain or share electrons during a reaction

state the maximum number of electrons the first shell of an atom can hold

2

state the maximum number of electrons the second shell of an atom can hold

8

state the maximum number of electrons the third shell of an atom can hold

8

state the maximum number of electrons the fourth shell of an atom can hold

20

explain how to predict the electronic configurations of the first 20 elements in the modern periodic table

• the notation of the electronic configuration of the first 20 elements in the periodic table adds up to the atomic number of the specific elements

• e.g. boron - atomic number (5) - electronic configuration (2,3)

reminder: how to calculate electronic configuration using carbon

• carbon has an atomic number of 6, meaning there are 6 protons and therefore 6 electrons

• shells fill from the first shell to the third

• 2 of the electrons will be in the full first shell

• 4 of the electrons will be in the incomplete second shell

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

• the number of shells in the electronic configuration of an element is the same as the period it is in

• e.g. elements in period 3 have 3 outer shells

• the number of electrons in the outermost shell is the same as the group it is in

• e.g. elements in group 7 have 7 elements in their outermost shell

describe the limitations of dot-and-cross diagrams in showing ionic bonding

• don’t show the relative sizes of ions

• don’t show ions are arranged

describe the limitations of ball-and-stick models in showing ionic bonding

suggests there are gaps between ions which is not true

describe the limitations of 2D representations in showing ionic bonding

doesn’t show the arrangement of ions in all layers

describe the limitations of 3D representations in showing ionic bonding

only shows the outer layer of ions

explain how ionic bonds are formed

• transfer of electrons

• between metal and non-metal atoms

• to produce cations and anions

recall what an ion is

• an atom or group of atoms

• with a positive or negative charge

state how to calculate number of protons in an atom given its atomic number and mass

number of protons = atomic number

state how to calculate number of electrons in an atom given its atomic number and mass

number of electrons = atomic number

state how to calculate number of neutrons in an atom given its atomic number and mass

number of neutrons = atomic mass - atomic number

explain the formation of ions in ionic compounds in groups 1 and 2

• atoms of elements in groups 1 and 2 will form cations

• as they will lose electrons

• due to these atoms having a lower number of electrons in their outermost shells

• meaning it is easier to complete their outer shell by losing electrons

explain the formation of ions in ionic compounds in groups 6 and 7

• atoms of elements in groups 6 and 7 will form anions

• as they will gain electrons

• due to these atoms having a higher number of electrons in their outermost shells

• meaning it is easier to complete their outer shell by gaining electrons

• resulting in an overall net negative charge

explain the use of the ending -ide in naming compounds

used when a compound contains only 2 atoms

explain the use of the ending -ate in naming compounds

used when a compound contains three or more atoms, one being oxygen

reminder: how to calculate the formulae of ionic compounds using the crossover rule (aluminium sulfate)

• aluminium ion - Al³⁺

• sulfate ion - SO₄²⁻

• aluminium sulfate compound - Al₂(SO₄)₃

explain the structure of an ionic compound

• giant lattice

• regular arrangement of ions

• held together between strong electrostatic forces of attraction

• between oppositely-charged ions

explain why ionic compounds have high boiling and melting points

• due to the large amount of energy needed

• to overcome

• strong electrostatic forces of attraction

• between oppositely-charged ions

explain why ionic compounds don’t conduct as solids

• ions are held in fixed positions

• so they cannot move

• and therefore cannot carry charge

explain why ionic compounds conduct when molten or in aqueous solutions

• ions are not held in fixed positions

• so they can move

• and therefore can carry charge

state the formula of aluminium

Al³⁺

state the formula of ammonium

NH⁴⁺

state the formula of bromide

Br⁻

state the formula of calcium

Ca²⁺

state the formula of carbonate

CO₃²⁻

state the formula of chloride

Cl⁻

state the formula of fluoride

Fl⁻

state the formula of sodium

Na⁺

state the formula of hydroxide

OH⁻

state the formula of oxide

O²⁻

state the formula of iodide

I⁻

state the formula of phosphate

PO₄³⁻

state the formula of lithium

Li⁺

state the formula of strontium

Sr²⁺

state the formula of potassium

K⁺

state the formula of sulfide

S²⁻

state the formula of nitrate

NO₃⁻

state the formula of sulfate

SO₄²⁻

explain how a covalent bond is formed

• a group of non-metal atoms

• sharing pairs of electrons

• to complete all atoms’ outer shells

recall the bonding used in the formation of molecules

covalent bonding

explain the formation of hydrogen as a simple covalent substance

• one single shared pair of electrons

• between two hydrogen atoms

explain the formation of hydrogen chloride as a simple covalent substance

• one single shared pair of electrons

• between one hydrogen atom

• and one chlorine atom

explain the formation of water as a simple covalent substance

• two single shared pairs of electrons

• between two hydrogen atoms

• and one oxygen atom

explain the formation of methane as a simple covalent substance

• four single shared pairs of electrons

• between four hydrogen atoms

• and one carbon atom

explain the formation of oxygen as a simple covalent substance

• one double shared pair of electrons

• between two oxygen atoms

explain the formation of carbon dioxide as a simple covalent substance

• two double shared pair of electrons

• between two oxygen atoms

• and one carbon atom

explain why simple covalent compounds have low melting/boiling points

• contain weak intermolecular forces between atoms

• which need low amounts of energy to overcome

explain why simple covalent compounds are poor conductors of electricity

• they contain no charged particles

• which therefore cannot move and carry charge

recall what type of covalent substance graphite is

giant covalent substance

recall what type of covalent substance diamond is

giant covalent substance

describe the structure of graphite

• each carbon atom is joined to three other carbon atoms by covalent bonding and one delocalised electron

• carbon atoms form hexagonal layered structure

• layers have weak intermolecular forces between them, allowing them to slide over each other

describe the structure of diamond

• each carbon atom is joined to four other carbon atoms by covalent bonding

• carbon atoms form a regular tetrahedral structure

• no free electrons

explain why graphite is used to make electrodes

• graphite contains a sea of delocalised electrons

• which are capable of moving and carrying charge

• allowing them to conduct electricity

• making it a good material for electrodes

explain why graphite is used as a lubricant

• graphite contains weak intermolecular forces between hexagonal layers

• allowing the layers to slide over each other

• making graphite a soft material and good to use as a lubricant

explain why diamond is used in cutting tools

• each carbon atom is joined to four other carbon atoms by covalent bonding

• giving diamond a rigid tetrahedal structure

• the rigidity allows diamond to be a good material for cutting tools

state the properties of buckminsterfullerene

• low melting point

• slippery

explain why buckminsterfullerene has a low melting point

• low amount of energy needed

• to overcome weak intermolecular forces

• between buckminsterfullerene compounds

explain why buckminsterfullerene is slippery

• weak intermolecular forces

• between buckminsterfullerene layers

• allowing the layers to slide over each other

state the formula of buckminsterfullerene

C₆₀

state the properties of graphene

• strong

• high melting point

• lightweight

• thermal and electrical conductor

explain why graphene is strong

strong covalent bonds between carbon atoms

explain why graphene is lightweight

graphene is a single layer of graphite

explain why graphene has a high melting point

• large amount of energy needed to overcome

• strong covalent bonds

• between carbon atoms

explain why graphene is a conductor

• contains a sea of delocalised electrons

• which are capable of moving and carrying charge