chemisty year 10 sem 1

element symbols

- represent the names of different elements

- consist of one or two letters

- first letter is always capital, if there is a second letter then it is lowercase

structure of periodic table (colums/rows)

- horizontal rows are called periods, and each one is given a number starting from the top

- vertical columns are called groups, and each one is given a number starting from the left

metals, non-metals and metalloids location on periodic table

metals are on the left side of the periodic table

non-metals are on the right side of the periodic table

metalloids are on a staircase shape between the metals and non-metals

metals

form 80%~ of the periodic table

group 1 (alkali metals): low melting points, soft, and very reactive

group 2 (alkaline earth metals): relatively soft and reactive

group 3-12 (transition metals): form coloured compunds, some are magnetic.

- only metals that are not silver in colour are copper (Cu) and gold (Au)

non-metals

while only 18 elements are considered to be non-metals, they make up most of the atomosphere and crust of Earth as well as living organisms

group 17 (halogens): only group to contains gases, a liquid (Bromine), and solids, most reactive group of non-metals

group 18 (noble gases): all gases and have a full valence shell, so are very unreactive

metalloids

- properties of both metals and non-metals

subatomic particles location in atom

protons and neutrons are tightly packed in the nucleus of an atom so they can only vibrate in position

electrons move in regions of space around the nucleus

cannot determine where an electron (we describe the region of space is electron is likely to be in) is and how fast it moving

region of space = electron cloud

properties of sub-atomic particles

number of protons, electrons, neutrons in atoms

neutral atoms - same number electrons and protons

atoms usually have the same number of neutrons and protons

adding a proton makes a new kind of atom

adding a neutron makes an isotope of that atom, a heavier version of that atom

atomic number and atomic mass

atomic number - number of protons in an element and number of electron in a neutral atom of the element

atomic mass - number of protons and neutrons added together in an isotope of an element

• mostly not whole numbers on the periodic table

• because masses are a weighted average of all the isotopes of that element

the bottom thing is called an atomic symbol

isotopes

atoms with the same number of protons but a different number of neutrons are called isotopes

named by their mass number as they have a different mass (because of the different number of neutrons they contain)

electron configuration

arrangement of electrons around the nucleus

• electrons occupy energy levels

• first energy level: max 2 electrons

• second energy level: max 8 electrons

• third energy level: 8 electrons (not max)

• fourth energy level: 2 electrons (not max)

an abbreviated way of writing this electron structure would be: 2, 8, 8, 2

this is called its electron configuration.

orbitals

electrons can only be in an energy level if they have the correct amount of energy

energy level can contain sub-levels called orbitals

orbital contains a maximum of two electrons

in energy level diagrams, electrons are drawn in pairs, showing they occupy the same orbital.

Only done if an energy level contains more than 4 electrons

electron configuration and the periodic table

element’s position can tell its electron configuration

period number determines number energy levels

• eg: period 2 elements have 2 energy levels and period 6 elements have 6 energy levels

group number determines the number of valence electrons

• eg: group 2 elements have 2 valence electrons and group 14 elements have 4 valence electrons

electron configuration + periodic table + valence electrons

valence electrons influence → element’s chemical properties

same group → similar properties → same number of valence electrons

excited electrons

in an atom, electrons are found in specific areas around the nucleus → electron shells

when thermal/electrical energy is added → electrons in the outermost shell jump to a shell further away from the nucleus

when an electron has jumped to a higher electron shell → excited electron

excited electrons eventually drop back to their original shell (extra energy is emitted as light/a photon)

colours of light emitted in excited electrons

colour of light emitted depends on the element

used to identify different elements

• eg: compounds with sodium → yellow colour

• eg: potassium → pale purple

electrons and chemical reactions (1)

in chemical reaction → valence electrons of atoms interact

electrons are shared/exchanged between atoms during the reaction

at the end of reaction:

• atoms either have filled their valence energy levels to 8 (non- metals)

• atoms empty their valence energy levels (metals)

eg: fluorine is a non-metal with 7 valence electrons, so it will gain 1 more to have a full valence energy level

• hydrogen can either empty it valency energy level or fill it up to 2 depending on the reaction

electrons and chemical reactions (2)

depending on which makes the least change to the atom, they either:

• gain electrons to fill their valence energy level to 8

• lose electrons to empty their valence energy level

atoms with a full/empty valence energy level are more stable (why noble gases are stable)

ions and valency

atoms which have gained/lost electrons are ions

they are no longer neutral (became charged)

the charge on an ion (eg: 2+, 3-) is called its valency

most atoms can only form one ion but some metals can have different valencies

when naming the ion; the valency is written in Roman numerals in brackets after the element name

• Fe²⁺ → Iron (II)

• Fe³⁺ → Iron (III)

positive ions

when atoms lose electrons

number of electrons lost → amount of positive charge

valency is written at the top right of the symbol

Al³⁺: an aluminium ion form when an aluminium atom loses 3 valence electrons

Li⁺: a lithium ion forms when a lithium atom loses a valence electron

Ca²⁺: a calcium ion forms when a calcium atom loses 2 valence electrons

negative ions

when atoms gain electrons

number of electrons gained → amount of negative charge

valency is written at the top right of the symbol

F^-: a fluoride ion forms when a  fluorine atom gains 1 valence electron

P^3-: a phosphide ion forms when a phosphorus atom gains 3 valence electrons

S^2-: a sulfide ion forms when a sulfur atom gains 2 valence electrons

ions and the periodic table

metal elements → positive ions

non-metal elements → negative ions

more than 4 = gains electrons

less than 4 valence electrons = loses electrons

except for metaloids (form covalent bonds) and some metals who lose or gain

(except hydrogen which forms either)

trends in periodic table - atomic size

atom’s size is measured using the radius of the atom

a new energy level is added with each period → atom grows bigger → atomic size increases from top to bottom

negative electrons are pulled in more tightly as the number of positive protons in the nucleus increases → atomic size decreases from left to right

trends in periodic table - reactivity METALS

metal reactivity increases from top to bottom and right to left

larger metal atoms have a weaker hold on their valence electrons → more readily removed → causing chemical reaction to occur

trends in periodic table - reactivity NON-METALS

metal reactivity increases from bottom to top and left to right

smaller non-metal atoms with more electrons in their valence shell → more capable of ripping electrons off other atoms to fill their valence shell

Compare sodium and magnesium in terms of their size and reactivity.

Bonding

In a chemical reaction, the valence electrons of atoms are shared or exchanged

At the end of the reaction, atoms have either filled or emptied their valence shell

Reactions cause atoms to join together, forming chemical bonds

Types of Bonding

Metallic bonding: metals only

Covalent bonding: non-metals only

Ionic bonding: metals and non-metals

Metallic Bonding

metal atoms have a weak hold on their valence electrons

valence elections are free to move through out the metal without being bound to one particular atom

the electrons are “shared” by the metal atoms are called delocalised atoms

eg: copper, iron, zinc (CIZ)

Properties of Metallic Bonds

• Metals are good conductors of electricity → electrons can move freely through the metal allowing them to carry an electric current

• This is why metals are also good conductors of heat- the electrons carry the energy through the substance

• Metals are malleable; can be bend or reshaped. Even if the metal atoms are moved around, the electrons will stay between them and hold them together

• Metals have high melting and boiling points → positive metal atoms are attracted to negative electrons → a lot of energy is needed to separate them

Covalent Bonding:

Non-metal atoms bond with other non-metal atoms by sharing their valence electrons

Electrons are shared so that the valence shell for each atom is ful

eg: water, hydrogen, carbon dioxide

Properties of Covalent Bonding

• Covalent substances do not conduct electricity or form ions (even when dissolved) and their electrons are trapped in the bonds between atoms → can’t carry an electric current

• Brittle and break easily → bonds within covalent molecules are very strong, the bonds between them are weak and easy to break

• Low melting and boiling point for the same reason

Ionic Bonding

Form when a metal atoms transfers its valence electrons to non-metal atoms

Metal becomes positively charged and non-metal becomes negatively charges

Ionic bond forms from the electrostatic attraction between the oppositely charged ions

Each ion is surrounded by ions of the opposite charge, building up a 3D structure called a lattice

eg: table salt; sodium chloride

Properties of Ionic Bonds

• Ionic solids cannot conduct electricity → ions hold each other in place so particles can’t move but when ionic substance is melted/dissolved → ions can move freely and conduct electricity

• Ionic compounds are brittle. Large enough force can push ions with the same charge close together → like charges repel → push each other apart

• Force of attraction between opposite ions in the lattice is very strong → high melting and boiling points

Forming of Ionic Compounds

• Positively charged ions (cations) bond with negatively charged ions (anions) through electrostatic forces of attractions (positive and negative charges attract)

• Bond to create a compound with a net (total) charge of zero

  • sum of neg charges = sum of positive charges

Polyatomic Ions

composed of more than one atom

create a compound with a net (total) charge of zero

more than 1 of a polyatomic ion → brackets around it
before writing the number in subscript

eg: Mg(NO₃)₂

Naming Ionic Compounds

1. name of the cation first

2. name of the anion with the ending changed to “ide“ second

3. Numbers are ignored.

Naming Ionic Formulae

1. symbol of cation first and symbol of anion second

2. If necessary, adjust the number of one or both ions so that the total positive charge equals the total negative charge

3. Remove charges and use subscripts to indicate the quantity of each ion. 1s are not written.

Electron Dot Diagrams

also called Lewis dot diagrams

Electron dot diagrams show the electrons in the valence energy level.

bc chemical reactions only involve atoms’ valence electrons

can be used to show bonds between elements