C1 science

• Subatomic particle Relative charge Location Relative mass

  Proton +1 nucleus 1

  Neutron 0 nucleus 1

  Electron -1 around nucleus 1/1840
• Electrons

  Have a small mass

  Move around the nucleus in electron shells

  Negatively charged

  Tiny but their shells cover up a lot of space

  Tiny mass

  Size of their shells determine the size of the atom
• Nucleus

  Middle of the atom

  Contains protons and neutrons

  Positive charge because of the protons

  Almost the whole mass is concentrated in the nucleus

  Is tiny compared to the size of the atom
• All matter is made up of tiny particles called atoms

  Atoms are tiny spheres that cannot be broken down

  The atoms in an element are all identical

  Atoms usually have no charge. Equal numbers of protons and electrons in an atom

  Most of the atom is empty space

  The properties of an element depend on its atomic structure. Atoms of the same element contain the same number of protons
• Dalton atomic theory

  In the 19th century John Dalton developed an atomic theory.

  He described atoms as solid spheres – different spheres made up all the different elements

  Atoms cannot be created or destroyed (broken down into anything simpler)

  Atoms are rearranged in a chemical reaction

  Compounds are formed when two or more different kinds of atom join together
• Thomson's atomic theory

  In 1897 JJ Thomson found out, through experiments, that atoms were not solid spheres

  He found, by measuring charge and mass in his experiments, that atoms contain smaller negatively charged particles – called electrons

  The plum pudding model showed the atom as a ball of positive charge with electrons stuck in it

  
• Rutherford's atomic theory

  In 1909 Rutherford and two of his students conducted the gold foil experiment. They fired positively charged alpha particles at an extremely thin sheet of gold

  When alpha particles came near the concentrated positively charged nucleus they were deflected. If they were fired directly at the nucleus they were deflected backwards. Otherwise, they passed through the empty space
• Rutherford's atomic theory

  From the plum pudding model they were expecting the particles to pass straight through the sheet or be slightly deflected at most. This was because the positive charge of each atom was thought to be very spread out through the atom. But whilst most of the particles did go straight through the gold sheet some were deflected more than expected and a small number deflected backwards. So the plum pudding model could not be right
• Rutherford's atomic theory

  He came up with the idea to explain the new evidence- the nuclear model of the atom

  In this there is a tiny positively charged nucleus at the centre where most of the mass is concentrated

  Negative electrons surrounds this nucleus so most of the atom is empty space
• Bohr model

  Electrons are contained in electron shells or fixed orbits

  That electrons orbit the nucleus in fixed shells and aren't anywhere in between. Each shell is a fixed distance from the nucleus and has a fixed energy

  
• Mass number is the formula of proton plus neutron (number at the top)

  Atomic number is the number of protons (number at the bottom)
• Periodic table

  Groups(vertical columns)- elements with similar properties and same number of electrons in the outer shell

  Period(row)- how many shells and have order of increasing atomic number
• First periodic table

  Created by Dimitri Mendeleev's

  Arranged fifty elements into the table of elements

  Began by sorting elements into groups based on their properties

  At first it was arranged in order of relative atomic mass but some elements did not fit the trend

  Left gaps to predict the properties of undiscovered elements
• Ions

  In an ion, which is an element that has lost or gained an electron, the number of protons is not equal to the number of electrons. Whereas in a neutral atom they are equal.

  Ions are charged particles they can be single atoms or groups of atoms

  When atoms lose or gain electrons to form ions all they're trying to do is get a full outer shell

  Negative ions (anions) form when atoms gain electrons- they have more electrons than protons. Negative charge gains an electron

  Positive ions (cations) form when atoms lose electrons- they have more protons than electrons
• Isotopes are different forms of the same element, which have the same number of protons but a different number of neutrons

  Relative atomic mass of an element is the average mass of one atom of the element compared to 1/12 of the mass of one atom of carbon-12

  The AR (Relative atomic mass) of an element is the average mass of the isotopes of the element so it is unlikely to be an exact whole number

  Relative atomic mass (AR) = Sum of (isotope abundance x isotope mass number)/ Sum of abundance of all the isotopes
• Ions

  Metals lose electrons and become positive ions

  Non-metals gain electrons and become negative ions
• Electron configuration

  Electrons are arranged in shells that orbit the nucleus. These shells are energy levels. The lowest energy levels are filled out first

  First shell – max of 2 electrons

  Second shell – max of 8 electrons

  Third shell – max of 8 electrons

  The number of shells which contains electrons is the same as the period of the element

  The group number tells you how many electrons occupy the outer shell of the element
• Ionic bonding

  The transfer of electrons. It is between metals (which lose electrons) and non-metals which gain electrons.

  The opposite charges cause them to become strongly attracted

  to each other by electrostatic forces.
• Ionic compounds

  Are made up positively charged ions and negatively charged ions

  Have a structure called a giant ionic lattice

  The ions form a closely packed regular lattice arrangement and there are very strong electrostatic forces of attraction between oppositely charged ions, in all directions in the lattice.

  They all have high melting point and high boiling points due to many strong bonds between the ions. It takes a lot of energy to overcome this attractions

  When they are solid the ions are held in place so the compound are free to move around
• Ionic compounds

  When ionic compounds melt the ions are free to move and they will carry electric charge

  Some ionic compounds dissolve in water. The ions separate and are all free to move in the solution, so they will carry electric charge
• When a ionic substance melts, the giant ionic lattice is broken.

  In its molten (melted) state, the ions are free to move around.

  To free the ions during melting, the ionic bonds must be broken.

  The ionic bonds (also called electrostatic forces) are strong, so lots of energy is needed to break them.

  Some ions have more than one charge so need more energy to overcome electrostatic forces e.g. MgO

  This means ionic compounds have high melting points
• For a compound to conduct electricity, it must have charged particles that are free to move

  When solid, the ions are locked into a lattice and cannot move.

  Therefore solid ionic substances do not conduct electricity.

  However, when molten or dissolved in water, the ionic lattice has been broken.

  Therefore the ions are free to move and carry charge.

  Molten or dissolved ionic substances do conduct electricity.
• Covalent bonding

  A covalent bond is a strong bond that forms when a pair of electrons is shared between two atoms

  Sharing of electrons

  Simple molecular substances are made up of molecules containing a few atoms joined together by covalent bonds.

  Two non-metals

  Valency number of covalent bonds formed by atoms
• Properties of metals- shiny, malleable, good conductor of electricity (because of delocalised electrons), ductile, high melting point and high boiling point

  In a metal the atoms lose many of their outer electrons

  The strong electrostatic forces means that metals have high melting and boiling points because these forces are strong a lot of energy is needed to break them
• As a solid, the metal atoms are packed close together in a regular

  structure. The outer energy level (shell) electrons are lost from each atoms and become free to move around throughout the metal. This

  leaves a giant structure (lattice) of positive metal ions surrounded by delocalised electrons.
• Metallic bonding

  The structure is held together by the electrostatic attraction between the positive metal ions and the negative delocalised electrons. This attraction is strong and metals have high melting and boiling points.

  They conduct electricity because of the amount of delocalised electrons. The electrons let electricity flow through them
• hydroxide ions OH−

  Nitrate ion NO−₃

  Carbonate ion CO2−3

  Sulfate ion SO2−4
• Polymers are made of covalently bonded carbon chains

  Polymers are molecules made up of long chains of covalently bonded carbon atoms. An example is polythene.

  They are formed when lots of small molecules called monomers join together
• Polymer for example polythene

  Type of bonding - strong covalent bonding within chains

  Intermolecular forces - many weak intermolecular forces of attraction along the length of the molecules

  Conduct electricity- poor conductor of electricity

  State at room temperature - solid
• Simper molecule for example water

  Size - small

  Type of bonding - strong covalent bonding within molecules

  Intermolecular forces - weak intermolecular forces between molecules

  Conductor electricity - poor conductor of electricity

  State at room temperature - liquid/gas at room temperature
• Intramolecular (covalent) forces between atoms are strong

  Intermolecular forces between molecules are weak and easily broken
• Giant covalent structures all the atoms are bonded to each other by strong covalent bonds

  They have very high melting and boiling points as lots of energy is needed to break the covalent bonds

  Generally don't contain charged particles, so they don't conduct electricity
• Diamond

  Rigid network of carbon atoms that each form four covalent bonds

  Strong covalent bonds take lots of energy to break, so diamond has a high melting point

  Strong covalent bonds also hold the atoms in a rigid lattice structure making diamond really hard

  Used in cutting tools

  Doesn't conduct electricity because it has no free electrons or ions
• Graphite

  Each carbon atom only forms three covalent bonds creating sheets of carbon atoms arranged in hexagons

  Aren't any covalent bonds between the layers they're only held together weakly, so they're free to move over each other. This makes graphite soft and slippery so its used as a lubricating material

  High melting points - the covalent bonds in the layers need loads of energy to break
• Graphite

  Only three out of each carbon's four outer electrons are used in bonds, so each carbon atom has one electron that's delocalised (free) and can move. So it conducts electricity and is used to make electrodes
• Graphene

  Is a type of fullerene

  One layer of graphite

  It's a sheet of carbon atoms joined together in hexagons

  The sheet is just one atom thick making it two dimensional substance

  Single sheet of carbon with free electrons conducts electricity well

  Light not heavy
• Buckminsterfullerene - C60

  Spherical fullerene made up of 60 carbon atoms

  Weak intermolecular forces between simple molecules

  Strong bonds within molecule but weak forces between molecules

  Soft and slippery

  Forms a hollow sphere made up of 20 hexagons and 12 pentagons

  Stable molecule that forms soft brownish-black crystals
• Metals have a giant structure

  The electrons in the outer shell of the metal atoms are delocalised (free to move around). There are strong forces of electrostatic attraction between the positive metal ions and shared negative electrons

  These forces of attraction hold the atoms together in a regular structure and are known as metallic bonding. Metallic bonding is very strong

  Compounds that are held together by metallic bonding include metallic elements and alloys

  It's delocalised electrons in the metallic bonds which produces all the properties of metals
• Empirical formula - the simplest whole number ratio of atoms of each element in a substance
• Formula (Mr) mass - the sum of the relative atomic mass (Ar) of all the atoms or 'ions' in a formula