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Flashcards in Topic 1 Deck (68):


All particles touch, they do not move/vibrate, have no/little energy



All particles touch but are moving, they flow to fit the space, have some energy



All particles aren't touching and are far apart, move freely, have high energy


Solid to Liquid



Liquid to Solid



Liquid to Gas



Gas to Liquid



Solid to Gas



Gas to Solid



Solid to Liquid process (Liquid to Solid is process but reversed)

Particles move faster as heat is added, begin to vibrate more, force of attraction weakens, can move around relatively freely, no longer in fixed arrangement, particles gain more energy


Liquid to Gas process (Gas to liquid is process but reversed)

Particles begin to move even faster, no arrangement at all, force of attraction becomes extremely weak, particles gain more energy, particles can move around complete freely



A substance is put in a solvent (substance that dissolves solute) to reduce its concentration



The movement of particles from an area of high concentration to an area of low concentration


Experiment 1 - Diffusion

- Hydrochloric acid placed at end of tube, ammonia solution at the other
- Where they meet, a ring of ammonium chloride appears
- Ammonia particles are lighter as they traveled further than the hydrochloric acid in the same amount of time


Experiment 2 - Dilution

- Put a coloured substance in a solvent (food colouring in water)
- Over time the colour levels out but is weaker than original colour - this is dilution
- Dilution at different temperature; more heat; more energy; more movement; so particles move to different areas quickly. So applying heat decreases time to fully dilute



Made up of a nucleus (protons and neutrons) and electrons orbiting around (in shells)



Two or more atoms bonded together



A group of the same type of atoms with the same number of protons. These appear on the periodic table.



A substance made up of two or more elements chemically bonded together.



A substance made up of two or more other substances not chemically bonded.



Method used to separate two substances, one solid, on liquid, that do not mix
- The mixture runs through a funnel with filter paper on it into a beaker or conical flask
- The solid is trapped in the filter paper
- The liquid runs into the container


Simple Distillation

Used to separate mixtures of a liquid and a soluble solid
- Liquid in the mixture boils in a round bottomed flash and move up, as it is a gas
- It enters the condenser - where cool water is flowing around a tube
- It condenses again
- Drips into the container at the end of the condenser
- Soluble solid left in around bottomed flask and liquid in beaker


Fractional Distillation

Used to separate two liquids with different boiling points
- Liquid Am the one with the lowest boiling points boils first and travels up the fractionating column
- Any of liquid B that evaporates condenses and drips back down, as the beads in the fractionating column are only at the temperature of liquid A (too cold)
- Liquid A travels through the condenser, condenses again, and drips out into a contrainer
- Liquid B remains in the round bottomed flask



Used to create pure crystals of a solid that has been dissolved in a solvent (liquid)
- Mixture is warmed
- Solvent evaporates
- Left to cool
- Solid forms crystals


Paper Chromatography

Used to test to see the different dyes in a substance (or to separate them)
- Paper with a pencil line and a dot of substance A is put into a beaker of solvent
- Paper must not be dipped in further than the pencil line
- The solvent will travel up the paper
- The dyes will (usually) separate, travelling up the paper
- The further away from the pencil line, the more soluble the dye (due to particle size)
- If substance A does not separate, it is either pure or not soluble in the solvent
- Sometimes, other known substances (B, C & D) will be placed alongside A to see which one is present


How chromatograms identify the composition of a mixture

Chromatography paper is placed in a solvent, the different compounds will travel at different speeds (due to the size of their particles)

RF value is calculated by distance moved by compounded divided by distanced moved by solvent
- Never more than one
- The higher the RF value, the higher the spot

The shape on the chromatogram can be compared with that of known substances and where they match, they are the same substance. Chromatograms are used to tell what substances make up a compound or mixture.



+1 Charge, Relative Mass - 1



0 Charge, Relative Mass - 1



-1 Charge, Nearly 0 or 1/1836


Atomic Number

# of protons (same as # of electrons)


Mass Number

# of protons + neutrons



Atoms of an element with different # of neutrons


Relative Atomic Mass (Ar)

The mass of one atom of an element


Calculation for RAM (Ar)

(% of isotope * its mass) + (% of isotope * its mass) / 100


Periodic Table Arrangement of Elements

Periodic table is in order of atomic number
Left and Right - Row is called a period
Up and Down - Column is called a group

Group # represents # of electrons on outer shell
Row # corresponds with # of shells


Calculate relative formula masses (Mr) from relative atomic masses (Ar)

Mr is relative formula mass, it is the mass of a molecule. Just add the Ar (relative atomic mass) of the atoms in the compound.



One mole of any atom or molecule of any substance will have a mass (in grams) that is equal to the Ar (or Mr) of that substance

One mole - 6.023 * 10^23


Mole Calculation using Ar and Mr

Ar / (Moles * Mr)


Balanced Equations

Balanced equations is when the reactants and products have the same number of each element


State Symbols

S - Solid
L - Liquid
G - Gas
Aq - Aqueous (solid dissolved in liquid)


Understand how the formulae of simple compounds can be obtained experimentally, including metal oxides, water and salts containing water of crystallisation

- Weigh the compound
- Remove one element of it through a reaction
- Weigh again
- If the first weight (mass) of the compound is AB and the second weight (mass) is A, you can work out B by: AB - A
- Work it out by diving the weight


Empirical Formula

The empirical formula of a compound gives the simplest whole number ratio of atoms of each element in a compound. If can be calculated from knowledge of the ratio of masses of each element in the compound.

A compound that contains 10g of hydrogen and 80g of oxygen has an empirical formula H20
Amount of hydrogen atoms = mass in grams/Ar of hydrogen (10/1) = 10 mol
Amount of oxygen atoms = mass in grams/Ar of oxygen (80/16) = 5 mol
Therefore formula/ratio is 10:5 or simplified to be 2:1


Reacting Masses

Typical Exam Question:
Calculate the mass of magnesium oxide that can be made by completely burning 6g of magnesium in oxygen

Equation for reaction: 2Mg + O2 -> 2MgO

1. Calculate the amount, in moles, of magnesium reacted
Ar of Mg = 24
Amount of Mg (6/24) = 0.25 Mol

2. Calculate the amount of MgO formed
The equation tells us that 2 mol of Mg for 2 mol of MgO, hence the amount of MgO formed is the same as the amount of Mg reacted.
Amount of MgO formed is 0.25 mol

3. Calculate the mass of the MgO formed
Mr of MgO + (24+16) = 40
Mass of MgO = (0.25 * 40) = 10g


Carry out mole calculations using volumes and molar concentrations

Moles / Volume = Concentration


Formation of ions by the gain or loss of electrons

An ion is any atom (or groups of atoms) that has a charge/is electrically charged. Ions are formed due to the loss or gain of electrons.

If an atom gains an electron, it becomes a negatively charged ion. Non-metals do this, they form anions. Elements group 5-7

If an atom loses an electron, it becomes a positively charged ion. Metals do this, they form cations. Elements group 1-3

Group 0&8 are noble gases and are inert + unreactive, will not form ions



Losing electrions



Gaining electrons


Positively charged common ions (google the chemical formula cause ceebs)

Potassium, Sodium, Lithium, Hydrogen, Magnesium (Mg2), Calcium (Ca2), Aluminium (Al3), Silver, Copper, Ammonium (NH4), Barium (Ba2), Zinc (Zn2), Copper 2, Lead 2, Iron 2


Negatively charged common ions (google the chemical formula cause ceebs)

Fluoride, Chloride, Bromide, Hydroxide, Iodide, Nitrate, Oxide, Sulfide, Sulfate, Carbonate


Deduce the charge of an ion from the electronic configuration of the atom from which the ion is formed

if an atom has four or less electrons in its outer shell, it will lose these electrons when it ionically bonds with another atom, If it has more, then it will receive electrons

Atoms that receive electrons become NEGATIVE
Atoms that "give" electrons because POSITIVE


Explain, using dot and cross diagrams, the formation of ionic compounds by electron transfer, limited to combinations of elements from Groups 1, 2, 3 and 5, 6, 7 (google diagram cause ceebs)

Atoms react, bonding with each other, in an attempt to fill their outer shell. Full outer shells mean that the atom is stable. Ionic bonding involves oppositely charged atoms being formed due to a gain/loss of electron.


Understand ionic bonding as a strong electrostatic attraction between oppositely charged ions

Ionic bonding happens between two ions: they are attracted to each other due to their opposite charges, so we say the ions have electrostatic attraction. This attraction bonds them together into an ionic compound.

How to work it out (Charge of ion from electronic configuration of atom from which the ion is formed):
- How many electrons are on the outer shell
- How many shells does it have
- How many electrons will it take to fill the outer shell
- Now see if it takes more transferring to lose electrons (go down a shell) or to gain electrons (fill the shell)
- Which ever one takes the least transferring will be the route taken
- If its lost, it will have a + charge of the number of electrons, to empty the shell
- If it gained, it will have a - charge of the number of electrons it gained to fill the shell


Understand that ionic compounds have high melting and boiling points because of strong electrostatic forces between oppositely charged ions

Atoms react, bonding with each other, in an attempt to fill their outer shell. Full outer shells mean that the atom is stable. Ionic bonding involves oppositely charged atoms being formed due to a gain/loss of electrons.


Covalent Bond

A bond formed between atoms by sharing a pair of electrons (one from each atom)


Understand covalent bonding as a strong attraction between the bonding pair of electrons and the nuclei of the atoms involved in the bond

Electrons, being shared by atoms in a covalent bond, are attracted to the nucleus of each atom in the bond. Remember that electrons are negative and protons in the nucleus are positive.


The formation of covalent compounds by electron sharing

Hydrogen (H2)
Hydrogen atoms only have 1 electron and they only need 1 more to complete their shell (first shell only need 2 electrons). To complete their shell, they form a covalent bond. Only one pair of electrons is shared between them, this molecule is known as H2 (hydrogen gas)

Chlorine (Cl2)
Much like hydrogen, chlorine atoms also only need one more electron.

Hydrogen Chloride (HCl)
As hydrogen and chlorine both only need one electron to complete their outer shell, they can bond with each other forming HCl.

Water (H2O)
Oxygen atoms have 6 electrons on the outer shell and therefore need 2 more electrons to complete their shell. However, hydrogen only needs one. This means that 1 oxygen atom must bond with 2 hydrogen atoms.

Methane (CH4)
carbon has 4 outer electrons, therefore it needs 4 more to complete its outer electron shell. Hydrogen only needs 1 more electron, so 4 hydrogen atoms bond with 1 carbon atom.

Ammonia (NH3)
NItrogen has 5 electrons on its outer shell, so it needs 3 more to complete its shell, so it bonds with 3 other hydrogen atoms.

Oxygen (O2)
In oxygen gas, one oxygen atom shares TWO pairs of electrons with another oxygen atom to complete its shell, this is known as a double covalent bond.

Nitrogen (N2)
Nitrogen atoms have 5 electrons on their outer shell, they need 3 more electrons. This means that two nitrogen atoms share THREE pairs of electrons to fill their outer shell, this is known as a triple covalent bond.

Carbon Dioxide (CO2)
In carbon dioxide, 2 oxygen atoms share two pairs of electrons with a carbon atom, this forms two double covalent bonds

Ethane (C2H6)
In ethane there are 2 carbon atoms and 6 hydrogen atoms. Each of the 6 hydrogen atoms share their only electron with one of the two carbon atoms (each carbon atom bonds with 3 hydrogen atoms), the 2 carbon atoms then share their last electron with each other.

Ethene (C2H4)
In ethene, there are 2 carbon atoms and 4 hydrogen atoms. The 4 hydrogen atoms each share their only electron with one of the two carbon atoms, the two carbon atoms then share their last two electrons with each other, forming a carbon-carbon double covalent bond.


Understand that substances with simple molecular structures are gases or liquids, or solids with low melting points

A simple molecule (one with only a few atoms) will have a low melting point.


Explain why substances with simple molecular structures have low melting and boiling points in terms of the relatively weak forces between the molecules

A substance with a simple molecular structure is one that contains only a few atoms in a molecule. The intermolecular forces (between the molecules) are weak, so it doesn't take much energy or heat to break them, meaning they will melt and boil under low temperatures, as even small amounts of heat energy are enough to break the bonds.


Explain the high melting and boiling points of substances with giant covalent structures in terms of the breaking of many strong covalent bonds

A giant covalent structure is one with many atoms bonded together. To melt or boil them, you are not separating intermolecular bonds (between molecules), you are separating intramolecular bonds that keep the molecule together. These bonds are strong covalent bonds which take a lot of energy to break., so a lot of heat energy is required before the bonds will break to boil or melt; meaning they have high melting and boiling points.


Understand that a metal can be described as a giant structure of positive ions surrounded by a sea of delocalised electrons

In a metal, atoms come together into a lattice, the electrons become detached from their atoms, delocalised (detach or remove from a particular place or location), making the atoms positive ions.


Explain the electrical conductivity and malleability of a metal in terms of its structure and bonding.

Metals have delocalised electrons, electrons carry electricity; so because there are free electrons charge can pass easily through a metal.

The structure of a metal is with rows of atoms on top of one another, in pure metals as all the atoms will be the same size, the layers can slide easily over one another making them easy to bend.


Electric current

Flow of electrons, although it could be a flow of ions (as they have a charge).


Understand why covalent compounds do not conduct electricity

In covalent compounds there are no electrons free to move, this means there can be no transfer of electricity through a covalent compound.


Understand why ionic compounds conduct electricity only when molten or in solution

When ionic compounds are molten or in solution, the positive and negative ions separate meaning that there are ions free to flow and can also conduct electricity.


Describe experiments to distinguish between electrolytes and nonelectrolytes

Set up an electric circuit with an LED and a break in the wire, put both ends of wire into a solution/molten substance. If the LED lights up then there is a current flowing, this will only be able to happen if the solution is conduction: so it must be an electrolyte. Conversely if the LED does not light up then there is no current flowing, and so the solution has not conducted electricity meaning it must be a nonelectrolyte.


Understand that electrolysis involves the formation of new substances when ionic compounds conduct electricity

In electrolysis ionic compounds conduct electricity. Positively charged ions move to one end, negatively to the other, these are then turned into atoms (by losing their charge) and so new substances are formed.


Describe experiments to investigate electrolysis, using inert electrodes, of molten salts such as lead(II) bromide and predict the products

NOTE: The cathode (negative) attracts Pb2+ ions as they are positive, the anode (positive) attracts Br- ions as they are negative.

- As soon as the lead(II) Bromide melts(becomes molten), the ions become free to move around, this movement enables the ions move allowing a charge to flow, meaning electrolysis can take place.

- The electrodes are made out of carbon - which is inert (unreactive).

- Connect the electrodes to a power source

- The positive lead (II) ions are attracted to the cathode, which is the negative electrode. When they get there, they gain 2 electrons each from the electrode. This forms lead atoms (they are no longer ions as they have no charge). These 'fall' to the bottom of the container as molten lead.

- Bromide ions (negative) are attracted to the positive anode. When they get there, the extra electron which makes the bromide ion negatively charged moves onto the anode, this loss of the extra electron turns each bromide ion into a bromine atom. These join in pairs (bond covalently) to form bromine molecules (which is gas).

The half equations...

At the cathode: Pb2+ + 2e -> Pb

At the anode: 2Br -> Br2 + 2e-


Write ionic half-equations representing the reactions at the electrodes during electrolysis

At the positive electrode, electrons will be lost; to show this we write the lost electrons as products:
2Br- -> 2Br + 2e-
Make sure the charges are equal on both sides :1- -> 1-

At the negative electrode, electrons will be gained so we write them as reactants:
2H+ + 2e- -> H2
And to make sure the charges are the same on both sides : 0 -> 0