chem - key concepts (1)

• John Dalton published his own three-part atomic theory in 1803: all substances are made of small particles (atoms) that cannot be created/divided/destroyed, atoms of the same element are exactly alike, and atoms of different elements are different.
• J.J Thomson used a cathode-ray tube to conduct an experiment identifying an error in Dalton’s atomic theory in 1897: atoms can be divided into smaller parts.
• Ernest Rutherford developed a new model which said that most of the atom’s mass is found in a region in the centre called the nucleus.
• O2: moles= 32 / (16 x 2) = 32/32 = 1
• Cu: moles = 127 / 63.5 = 2
• Making the overall balanced equation – 2Cu + O2 -> 2CuO
• Therefore, you have a ratio of 2:1:2 for Cu: O2: CuO
• CuO: moles = 159 / (16 + 63.5) = 2
• An atom has a nucleus (contains protons and neutrons), surrounded by electrons shells.
• Neutron has charge = 0 and mass = 1, Proton has charge = +1 and mass = 1, Electron has charge = -1 and mass = 1/1836.
• Atoms contain equal number of protons and electrons because atoms are neutral and the charges on a proton are +1 and on an electron are -1 so the charges cancel each other out (ends up neutral).
• The nucleus of an atom is very small compared to the overall size of the atom but most of the mass of an atom is concentrated in the nucleus.
• Atoms of an element have the same number of protons in the nucleus and this number is unique to that element.
• Isotopes are different atoms of the same element containing the same number of protons but different numbers of neutrons (therefore having different masses).
• The relative atomic mass is calculated using the abundance of different isotopes and because it is an average it can lead to the ram not being a whole number.
• Elements with properties predicted by Mendeleev were later discovered and the gaps were then filled.
• Knowledge of isotopes made it possible to explain why the order based on atomic weights was not always correct, because some elements have a higher mass than others when isotopes are considered, but a lower one if you only look at one specific isotope.
• Elements are arranged in order of atomic (proton) number and so that elements with similar properties are in columns (groups).
• Elements in the same group have the same number of electrons in their outer shell, which gives them similar chemical properties.
• Metals are elements that react to form positive ions.
• Majority of elements are metals.
• The concentration of a solution can be measured in mass per given volume of solution.
• The moles of a substance can be calculated by dividing the mass by the molar mass of the substance.
• The molar mass of a substance is numerically equal to its relative formula mass.
• Stoichiometry refers to the balancing numbers in front of compounds/elements in equations.
• In a reaction with 2+ reactants, one is used in excess to ensure that all the other reactant is used, and this reactant is called the limiting reactant.
• Chemical reactions can be represented by symbol equations, which are balanced in terms of the numbers of atoms of each element involved on both sides of the equation.
• 127g Cu reacts with 32g of oxygen and 159g of CuO is formed.
• The number of particles in a mole of a given substance is the Avogadro constant: 6.02 x 1023.
• Balancing numbers in a symbol equation can be calculated from the masses of reactants/products: convert the masses in grams to amounts in moles (moles = mass/Mr) and convert the numbers of moles to simple whole number ratios.
• An example of a balanced equation is Cu + O2 -> CuO.
• Chemical amounts are measured in moles.
• The empirical formula of a compound can be determined by heating the compound to burning in a crucible and weighing the mass of the resulting magnesium oxide.
• The total mass is unchanged in a reaction in an open flask that takes in or gives out a gas, as some mass is lost when the gas is given off.
• One mole of particles of a substance is defined as the Avogadro constant number of particles (6.02 x 1023 atoms, molecules, formulae or ions) of that substance and a mass of ‘relative particle mass’ g.
• The molecular formula of a compound can be calculated by multiplying the empirical formula by the molar mass of the compound.
• The symbol for the unit mole is mol.
• The law of conservation of mass states that no atoms are lost or made during a chemical reaction, so the total mass of the products is equal to the total mass of the reactants.
• The mass of one mole of a substance in grams is numerically equal to its relative formula mass.
• The stoichiometry of a reaction can be deduced from the masses of the reactants and products.