Chemistry Notes - Atomic Structure PDF

The structure of atom Atom  The smallest particle of a chemical element that can exist Element  An element is a substance consisting atoms which all have the same number of protons - i.e. the same atomic number Proton  A stable subatomic particle...

The structure of atom Atom  The smallest particle of a chemical element that can exist Element  An element is a substance consisting atoms which all have the same number of protons - i.e. the same atomic number Proton  A stable subatomic particle occurring in all atomic nuclei, with a positive electric charge equal in magnitude to that of an electron Neutron  A subatomic particle of about the same mass as a proton but without an electric charge, present in all atomic nuclei except those of ordinary hydrogen Electron  A stable subatomic particle with a charge of negative electricity, has a relative mass 1/2000th that of proton Nucleus  The positively charged central core of an atom, consisting of protons and neutrons and containing nearly all its mass Atomic number  The number of protons in the nucleus of an atom, which is characteristic of a chemical element and determines its place in the periodic table Isotope  Forms of the same element with the same number of protons, but different numbers of neutrons Hydrogen has 3 isotopes each one has a different relative isotopic mass Mass number  The total number of protons and neutrons in a nucleus. Different isotopes of the same element have different mass numbers Relative atomic mass (Ar)  The ratio of the average mean mass of one atom of an element The structure of atom 1 The structure of atom 2 The structure of atom 3  Electronic arrangement Shells and quantum numbers The word “shellˮ actually refers to a particular energy level The shell which an electron is found in is described by the principal quantum number, n. Different shells have different quantum numbers. Electronic arrangement 1  Energy levels Shells can go from 2, 8 then 8 or 18. The word ‘shellʼ actually refers to a particular energy level. The shell a electron is in is described by the principle quantum number, n. Different shells have different quantum numbers. Orbitals and Shells Electrons within each shell donʼt have the same amount of energy so the energy levels or shells are broken down into subshells called orbitals Each shell can hold a different number of orbitals. Shells can be divided further into subshells (a group of the same type of orbitals within a shell) As the principal quantum number increases new types of orbital are contained within the shell. Shell 1 contains 1 s orbital Shell 2 upwards contain 3 p orbitals (as well as the s orbital) Shell 3 upwards contain 5 d orbitals (as well as the s and p orbitals) Electronic arrangement 2  Shell 4 upwards contain 7 f orbitals (as well as the s, p and d orbitals) Types of orbital There are four different types of orbital: s, p, d and f. An S-orbital is spherical in shape. P-orbitals are shaped like and hourglass. They can lie along the x, y or z axis. We call them Px, Py and Pz. An orbital is a region within an atom that can hold up to two electrons with opposite spins Electronic arrangement 3  Electrons are added one shell at a time Lowest energy level is filled first Each energy level must be full before filling the next level Each orbital is filled singly before pairing Electronic arrangement 4  3Li = 1S², 2S¹ 4Be = 1S², 2S² 5B = 1S², 2S², 2P¹ 6C = 1S², 2S², 2P² 7N = 1S², 2S², 2P³ 8N = 1S², 2S², 2P⁴ 9F = 1S², 2S², 2P⁵ 10A = 1S, 2S², 2P⁶ The Aufbau principle Niels Bohr developed the Aufbau principle which states that the lowest energy sub-levels are occupied first. This means the 1s sub-level is filled first, followed by 2s, 2p, 3s and 3p. However, the 4s sub-level is lower in energy than the 3d, so this will fill first. Electronic configurations  Fill the orbitals in order of increasing energy  Each orbital can hold a maximum of two electrons Electronic arrangement 5  Elements in the same group have the same number of electrons in their outer shell  Number of electrons = group numbers The Pauli exclusion principle and spin The Pauli exclusion principle states that each orbital may contain no more than two electrons. It also introduces a property of electrons called spin, which has two states: ‘up' and 'down'. The spins of electrons in the same orbital must be opposite, i.e. one 'up' and one 'down'. A spin diagram shows how the orbitals are filled. Orbitals are represented by squares, Rules for filling electrons When two electrons occupy a p sub-level, completely fill the same p orbital or half fill two different p orbitals. Hund's rule states that single electrons occupy all empty orbitals within a sub- level before they start to form pairs in orbitals. If two electrons enter the same orbital there is repulsion between them due to their negative charges. The most stable configuration is with single electrons in different orbitals. Electron configuration of Cr and C The 4s orbital contains one electron. This is because the 4s and 3d sub-levels lie very close together in energy, and the 3d being either half full or completely full is a lower energy arrangement. With larger atoms like this it can be useful to shorten the electron arrangement. Cr  1s² 2s² 2p⁶ 3s² 3p⁶ 3d⁵ 4s¹ Electronic arrangement 6  Ionic bonding Bonding  Atoms of elements combine in order to achieve full outer shells by: transferring or sharing electrons Ionic bonding This results from the electrostatic attraction between oppositely charged ions Metal atoms lose electrons to form positive ions Non-metal gain electrons to form negative ions The ions are arranged in a crystal lattice Molecular Ions Ionic bonding 1  Some complex ions are made up of more than one element. These are called compound or molecular ions. There are four in particular that you need to know the formulae and charges of: Nitrate NO3 sulphate SO4₂- carbonate CO3₂- Ammonium NH4 Charge on the ions Metals lose electrons to form positive ions while non-metals gain electrons to form negative ions. The number of electrons gained or lost by an atom is related to the group in which the element is found. The elements in groups 4 and 8 (also called group 0 do not gain or lose electrons to form ionic compounds. Ionic compounds Ionic bonding 2  Ionic bonds are attractions between positive and negative ions which are held together in a giant lattice. These attractions are very strong. Electrical conductivity Ionic compounds only conduct electricity when their ions are free to move This means when they are molten or in solution. To conduct electricity there must be a movement of charge Solubility Ionic compounds dissolve in polar solvents, such as water. The water molecules attract the positive and negative ions. The giant ionic lattice breaks down and water molecules surround the ions, forming a solution. Ionic bonding 3  Covalent bonding Covalent bonding usually takes place between non-metal atoms A single covalent bond consists of a shared pair of electrons The nucleus of each atom attracts the bonding electron of the atom to which it is bonded Properties of covalent bonds Covalent bonds are directional, this means they only attract the atoms involved in the bond Ionic bonding attracts in all directions Dative covalent bond (Co-ordinate bond) This is a type of covalent bond where both of the shared electrons has come from the same atom. An example of this is the bond between H and NH3, which produces NH4 Ammonia is able to do this because it has a pair of electrons that are not involved in a bond (lone pair of electrons) Covalent bonding 1  Dative covalent bonds are shown as arrows pointing towards the atom that the electrons are being shared with. Bond lengths and bond strengths Bond breaking is an endothermic process (you need to add energy to break the electrostatic attraction between the two atoms) Bond enthalpy is the energy needed to break a bond This will depend on the atoms involved and the number of bonds between the atoms The more bonds there are between the atoms, the stronger the bond and the more energy is needed to break the bond. The more bonds between atoms, the more electrons are being shared. This increases the attraction between the nuclei and the electrons, pulls the atoms closer together and makes the bond shorter Properties of simple covalent substances Low melting and boiling points At room temperature and atmospheric pressure they may exist as either solid, liquid or gas Donʼt conduct electricity, either as a solid or liquid Covalent bonding 2  Metallic bonding Metallic bonding is due to the attraction of a lattice of positive metal ions to a “seaˮ of mobile electrons. Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons Delocalised electrons are shared between more than two atoms A giant metallic lattice is a 3d structure of positive ions and delocalised electrons, bonded together by strong metallic bonds Properties of metals High melting and boiling points Attraction between positive ions and delocalised electrons is strong High temperatures needed to break the metallic bonds Metallic bonding 1  Good electrical conductivity Delocalised electrons are free to move anywhere in the lattice Malleability and ductility Because delocalised electrons can mov this allows some degree of give. Layers can slide past each other Metallic bonding 2  Mole calculations Relative atomic mass Relative molecular mass is the mean mass of a molecule compared to the mass a 12C atom Isotopes and relative atomic mass Calculating relative atomic mass (Ar) Calculating the relative atomic mass of an element given the percentage abundance of its isotopes Ar of Element = (mass x abundance) + (mass x abundance) / total abundance A sample of chlorine contains 80% of chlorine-35 and 20% chlorine-37. Calculate relative atomic mass of this sample of chlorine. 80  35  20  37  3540 3540/100  35.4  NaCl  2335.558.5  HNO3  11448163  63  Na2CO3  232 461248 163 106 Mole calculations 1  Moles Calculating moles: moles(mol)  Mass (g) / Mr (g/mol) Avogadroʼs number  Shows the number of particles in one of mole, its numerical value is 6.02  10^23 Mole calculations 2  Often only soluble in non-polar solvents Often form amorphous solids (no definite regular pattern to the particles in the solid) Giant covalent substances The atoms are held together by strong covalent bonds The whole crystal of a giant molecule is considered to be a single molecule For a giant molecule to melt each atom needs to be able to move freely, so the covalent bonds between the atoms need to be broken Examples of giant covalent molecules are carbon and silicon Properties of giant covalent substances Thousands of atoms bound together by covalent bonds High melting and boiling points Donʼt conduct electricity (except some allotropes of carbon) Insoluble in polar and nonpolar solvents Examples are carbon and silicon Carbon allotropes Diamond Very hard and least compressible (used as a cutting tool) Good conductor of heat at room temperature (can be used as a heat sink) Transparent when pure Each carbon is bonded to four other carbons Tetrahedral arrangement of the atoms in the crystal Does not conduct electricity Covalent bonding 3  Graphite Each carbon is bonded to three other carbons with bond angles of 120°. Atoms are arranged in layers. The unbonded electron in carbon becomes delocalised and can be found between the layers. The delocalised electrons allow graphite to conduct electricity. Weak forces of attraction between layers allow the layers to slide past each other (lubricants and pencil lead). Polar covalent bonding in some covalent bonds the pair of electrons is not shared equally. This is because the nucleus of one atom attracts the bonding pair of electrons more than the nucleus of the other atom. This means that the bonding pair of electrons spends more time with one atom, making it slightly negative. The other atom becomes slightly positive. The power of an atom in a molecule to attract the bonding electrons to itself is called electronegativity The greater the difference in electronegativity between the two atoms involved in the bond, the more polar the molecule Properties of polar covalent bonding Have higher melting and boiling points than non-polar substances in similar size Soluble in polar solvents like water Do not conduct electricity Covalent bonding 4  Shapes of molecules Valence shell electron pair repulsion (VSEPR) Electrons involved in a bond are called bonding electrons Bonding electrons will arrange themselves as far as possible from each other Electrons not involved in a bond are called lone pair electrons (non-bonding electron pair) Lone pair electrons repel the bonding pairs of electrons pushing them closer together Organic molecules Carbon has four valence electrons, so needs to form four bonds to achieve a full outer shell Carbon with four single bonds forms a tetrahedral shape You need to be able to draw the three-dimensional shape Covalent bonding 5  Polar & Non-polar molecules Polar covalent Non-polar molecule  A molecule where the electrons are distributed evenly throughout the molecule O2, H2, CO2 Polar molecular  A molecule with partial positive charge in one part of the molecule and similar negative charge to uneven electron distribution. H2O, NH3, CO Electronegativity The tendency of an atom in a molecule to attract the bonding electrons to itself is called electronegativity. The greater difference in electronegativity between the two atoms involved in the bond, the more polar in the molecule. Metals have a low electronegativity Polar & Non-polar molecules 1 Polar & Non-polar molecules 2  H2O  Oxygen) 3.44 minus 2.20 Hydrogen)  1.24 CO2  (oxygen) 3.44  (carbon) 2.55  0.89 Polar compounds Pair of electrons in a covalent bond isnʼt shared equally sometimes because the nucleus of one atom attracts the bonding pair of electrons more than the nucleus of the other atom. This means that the bonding pair of electrons spends more time with one atom making it slightly negative. The other atom becomes slightly positive Polar & Non-polar molecules 3  Intermolecular forces  The attraction or repulsion between neighbouring molecules. Dipole - separation of change within a covalent molecule Van Der Waals forces - all intermolecular attractions are Van Der Waals forces. London dispersion forces Also called temporary dipole - induced dipole forces They are weak forces present between non-polar covalent molecules. They are less than 1% of the force of a covalent bond more electrons → more movement → bigger dipoles → stronger attraction Dipole-dipole forces Another form of Van Der Waals forces These are permanent forces between polar molecules Ion dipole interactions Interactions that occur between ions and polar molecules Polar & Non-polar molecules 4  An example is when an ionic salt is dissolved in water Polar & Non-polar molecules 5  Empirical formula & percent yield  Calculate RFM for KMnO4 K  39 , Mn= 55 , O4  164 5539164  158  Moles in 2.842g of sodium sulfate, Na2SO4 Na2  232  46  32  64  142 Moles = mass / mr Moles  2.842 / 142 = 0.0197  6.02g of magnesium, number of moles in magnesium sulfate  MgSO4 Moles = mass / mr Moles  6.02 / 120 24  32  64  120 = 0.0501  80% of chlorine-35, 20% chlorine-37. Calculate RFM of chlorine. 80  35  2800 20  37  740 Empirical formula & percent yield 1  2800  740  3540 / 100  35.4 Empirical formula shows the relative numbers of atoms of each element present, using the smallest whole numbers of atoms The empirical formula of a compound is the simplest whole number ratio of ions in a formula unit or atoms of each element in a molecule How to find the empirical formula 1  Find the mole of each atom 2  Divide all the moles by the smallest number to find at least one whole number Empirical formula & percent yield 2  3  If still not whole numbers, multiply them all by a number to get whole numbers Examples A compound was found to contain 828 g lead and 64 g oxygen. Find its empirical formula. Moles  Mass / Mr 828 / 207.2  3.996  rounded to 4 64 / 16  4 Pb ₁0₁ A compound was found to consist of 9.6 g copper, 1.8 g carbon and 7.2 g oxygen. Find its empirical formula. copper 9.6 / 64  0.15 / 0.15  1 carbon 1.8 / 12  0.15 / 0/15  1 oxygen 7.2 / 16  0.45 / 0.15  3  Cu₁C₁O₃ Empirical formula & percent yield 3  A compound has a relative formula mass of 32. It consists of 87.5% nitrogen and 12.5% hydrogen. Find its empirical formula. RFM of N  14, H  1 87.5% N  87.5g 12.5% H  12.5g Concentrating & Balancing reactions Soluble - the substance that dissolves in a solvent to form a solution Saturated solution - a solution in which the maximum amount of solute has been dissolved Supersaturation - the difference between the actual concentration and the solubility concentration at a given temperature Reactions Many reactions take place in liquid solutions This allows for the reacting particles to move, collide and react Empirical formula & percent yield 4  A solution contains a solute dissolved in a solvent. Where the solution is a liquid, the solvent is a liquid. Water and ethanol are common solvents Solutes can be solid, liquid or gases. For example, carbon dioxide gas in a fizzy drink is the solute Example What is the molarity of a solution of NaOH if there are 4 moles of NaOH dissolved water to make 1 liter of solution? 4 moles / 1 liter = 4M Empirical formula & percent yield 5  0.050  solution 1.975  potassium 1.975 / 0.050  39.5g dm3 Balancing reactions Chemical equation  Describes a chemical change. Parts of an equation: Reactant  The chemical(s) you start with before the reaction Written on left side of equation Product  The new chemical(s) formed by the reaction Right side of equation Subscripts and Coefficients Subscript - shows how many atoms of an element are in a molecule. As an example  H2O  2 atoms of hydrogen H, 1 atom of oxygen O Coefficient - shows how many molecules there are of a particular chemical. As an example: 3 H2O  Means there are 3 water molecules. Empirical formula & percent yield 6  Law of conservation of mass In a chemical reaction, matter is neither created nor destroyed. In other words, the number and type of atoms going INTO a reaction must be the same as the number and type of atoms coming OUT. If an equation obeys the Law of Conservation, it is balanced. An Unbalanced Equation CH4  O2  CO2  H2O A Balanced Equation CH4  2O2  CO2  2H2O Empirical formula & percent yield 7  Rules  Matter cannot be created or destroyed.  Subscripts cannot be added, removed, or changed.  You can only change coefficients.  Coefficients can only go in front of chemical formulas...NEVER in the middle of a formula. Balance the following equation by adjusting coefficients N2  H2  NH3 Reactant Product N 2 1 (x 2) 2 H 2 (x3)  6 3 (x2)  6 N2  H2  NH3  N2  3H2  2NH Empirical formula & percent yield 8  Calculating percent yield Percent Yield Percentage yield is a way of comparing amount of product made (actual yield) to amount expected (predicted yield). Actual yield is the amount of product experimentally obtained by a chemical reaction. Percent yield = actual yield / theoretical yield x 100 Actual yield is measured using a scale or indirectly. It is less than theoretical Percent yield is the ratio of actual yield to theoretical yield multiplied, by 100% Step 1 Calculate the number of moles of the reactant. Step 2 Use the stoichiometry to calculate the theoretical number of moles of product. This is the theoretical yield) Step 3 In the question there will be the actual mass of the product. Calculate the actual number of moles of product (this is the actual yield). Calculating percent yield 1  Limiting reagent  Limiting reagent is the reactant that is completely consumed in a chemical reaction Calculating percent yield 2