General chemistry

Element = substance made from atoms with the same number of protons in the nucleus

Atom = smallest particle of an element that still has chemical properties

Molecule = 2+ atoms chemically joined

Compound = 2+ elements chemically joined

Mixture = 2+ elements not chemically joined

Ion = charged particle formed when an atom or group of atoms loses or gains electrons

Metal atoms generally lose electrons to form positive ions

Non-metal atoms generally gain electrons to form negative ions.

REACTANTS undergo a chemical change to create a PRODUCT.

Balanced equations are when there are the same number of each type of element on both sides of the equation.

When balancing word or symbol equations, you may only add numbers infront of the whole element or compound.

e.g. H2 + Cl2 -> 2HCl

A big number infront multiplies ALL elements in the compound if it is at the start of a compound such as HCl.

A little number after an element only multiplies that element, none before or after it.

State symbols include g- gas, s- solid, l- liquid and aq- aqueous solution. 

You can show the charge on an ion by placing 'n+' or 'n-' after it. 'n' is the charge number.

E.g. O2- 

or   Mg2+

Oppositely charged ions in a solution may join to form an insoluable solid- a precipitate.

The total electrical charge is the same on both sides of the equation.

Cation = electron with positive charge (usually form from hydrogens or metals)

Anion = Negatively charged ion. (Usually formed from non-metals)

If an atom does not have a full outer shell of electrons (2 in the first shell, 8 in the second, 8 in the fourth, 2 in the 5th) it is unstable. In order to stabalise it must lose or gain electrons so that it has a full outer shell. As in an atom as seen on the periodic table there are equal numbers of positive protons and negative electrons in an atom (giving it a neutral charge), if the atom gains electrons it will have an overall negative charge, but if it loses electrons (so has more protons than electrons) it will have an overall positive charge.

hazard = something that could cause someone to be in danger, such as water on the floor which could cause someone to slip.

We often use symbols on products to tell us what danger products my put us in.

Precations = something we do to reduce the risk from hazards, such as wearing goggles.

We now know that atoms have protons and neutrons in the nucleus, and electrons in shells around it. Here as some different historic models of the atom:

1803- Dalton- SOLID ATOM MODEL, he believed that all atoms of an element were identical.

1897/ 1904- Thompson- Discovers electrons/ PLUM PUDDING MODEL, where atoms are spheres of positive charge with negatively charged electrons inside.

1911- Rutherford - SOLAR SYSTEM MODEL, where atoms have a positive nucleus surrounded by negative electrons.

1913- Bohr- ELECTRON SHELL MODEL Electrons have shells/ energy levels around the nucleaus.

1918- Rutherford- Discovers the proton

1932- Chadwick- Discovers the neutron

Isotope = Versions of an element with the same number of protons (having a different number of protons would make it a different element) but a different number of neutrons in the nucleus.

- same atomic number

- different atomic mass

Relative atomic mass (Ar) is not the same as mass number.  Ar is the MEAN MASS of an atom of a certain element, compared to  1/12 of the mass of a Carbon 12 atom (tsking into account relative abundance of each isotope in a sample of an element. May not be a whole number.

• This version of table developed into modern periodic table. 

1. He put elements in order of Ar

2. Checked the properties of the elements and compounds (such as H2)

3. Swapped places of some elements so they would fit with other elements with similar/ the same properties.

4. Left gaps where he believed other elements would go. (these were later discovered and fit Mendeleeve's predictions very precisely)

Different groups have different properties. For example, Group 0 elements (which meneleeve did not predict as they are unreactive) are unreactive, Group 6 elements do not react with water, do react with oxygen, and have 6 electrons in their outer shell, and Group 7 elements All react with water, don't react with oxygen and have 7 electrons in their outer shell.

• Positively charged ions take the name of the element, e.g. H+ = Hydrogen
• Negatively charged ions formed froma single non-metal atom takes the name of the element with 'ide' on the end. E.g. F- = Fluoride
• Negatively charged ions in compounds with 3+ elements (one of which must be oxygen) end in 'ate'. E.g. (NO3)-   (one nitrogen atom, 3 oocygen atoms with an overall charge of -1)= nitrate
• Ionic bonds are electrostatic forces of attraction between oppositely charged ions.
• Ions in ionic compounds have lattice structure which has a regular arrangement of ions.
• Ionic compounds generally have high melting and boiling points.
• Ionic compounds are often soluable in water 9dissolve into aqueous solutions)
• For any ionic compound to conduct electricity, its electrons must be free to move around, e.g., when it is in molten state or dissolved in water (as then they can carry the electric charge between places)

Covalent bond = when a pair of electrons is shared ebtween 2 atoms.

• Strong
• Between non-metal atoms
• Often produce molecules

Models:

• Dot an cross diagrams with shells
• Dot and cross diagram without shells
• Structural formula

Bonding:

• Few atoms joined by strong covalent bonds
• e.g. H2
• e.g. H20
• low melting and boiling points
• Usually gas or liquid at room temperature
• Wax is an example of a simple molecular structure that is solid at room temperature
• Don't conduct electricity as not electrically charged or contain electrons that are free to move
• Often insoluable in water
• only dissolve in water if they can form strong enough intermolecular forces with water molecules.

• Contain many atoms
• Joine dby strong covalent bonds
• Lattice structure
• non metal elements or compounds
• generally have high melting and boiling points
• Solid at room temperature
• Lots of energy needed to break the bonds allowing them to melt.
• Insoluable in water

Models:

• Ball and stick

Diamond- each atoms bonded to 4 others, Strong covalent bonds between atoms.

Graphite- each atom bonded to 3 others, weak intermolecular forces between layers, strong covalents bonds between atoms in a layer.

Grapene- Each atom covalently bonded to 3 other atoms, regular lattive structure, resembles single layer of graphite

Fullerenes- Resemble sheet of Graphiene rolled into a tube or a ball:

Buckminsterfullerene ('buckyballs')- Carbon atoms arranged in pentagons or hexagons, conduct electricity due to having delocalised electrons, soften when in solid state due to having weak intermolecular forces.

nanatubes- hollow tube structure, closed or open ends, up to several mm long, conduct electricity due to having delocalised electrons, strong due to having strong covalent bonds.

• Most elements are emtal (left hand side of periodic table)
• Usually: shiny, good conductors, high density, high melting and boiling points.
• Non metals have opposite characteristics.
• Metals are malleable (If force is applied, layers of ions slide over each other and the metal shapes without shattering), non metals are brittle.
• Giant lattive of positively charged ions
• Sea of delocalised electrons (reason for conductivity)
• Metallic bonds are strong electrostatic forces of attraction between positive ions and delocalised electrons.
• Insoluable in water, though do occaisionally appear to dissolve as the react with water producing metal hydroxides which dissolve, along with Hydrogen.

Written formulae: Doesnt always show how molecules are structured, though does show how many of each atom and May show lowest ration (em[irical formula) . Structural formulae does give idea of arrangement.

Drawn diagrams: Don't give impression of space (2 dimentional) or show bonding and non-bonding electrons.

Ball and stick: Show how each atom is bonded and are 3d, but do not show bonding/ non-bonding electrons

Space-filling models: As with ball and stick models but show relative sizes but not types of bonds and all atoms may not be visable in complex molecules.

etc

Relative formula mass (Mr) = relative atomic masses of all atoms in the formula

HAS NO UNITS and only applies to covalent substances.

Empirical formula = simplest WHOLE NUMBER ratio of each type of atom in a compound.

To calculate:

Divide the given mass of each element by its atomic mass

DIvide each of these by the smallest of your answers (will have one for each element).

If fractions are involved, multiply all elements to make sure that all values are integers.

These are the numbers of each element that you will need.

Total mass of reactants is always the same as the total mass of products in a reaction. HOWEVER, this may not appear to be true if a gas is formed in a reaction taking place in a on-enclosed system (e.g. a container without a lid). THis is because the gas atoms move freely so will not stay in the container.

In addition:

• THe mass of rective metals may increase due to reacting with oxygen (creating a metal oxide0
• Mass of metal carbonate decreases if heated as CO2 is produced.
• In precipitation reactions, two souable reactants form an insoluable product. (precipitate)

A reactant is in excess if this is enough to react with all of the other reactant or there is some left over when the reaction stops.

A reactant is limiting if it runs out before another reactant, meaning that it stops the reaction.

Mass of product formed is controlled by the mass of limiting reactant.

Stoichiometry = ratio of the amounts of reactants and products formed.

https://www.youtube.com/watch?v=nZOVR8EMwRU