Chemistry Unit 1 - Periodicity and properties of elements Flashcards
what is an atom?
-the smallest particle of a chemical element that can exist
what is an element?
-a substance consisting of atoms which all have the same number of protons i.e. the same atomic number
what is meant by 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
what are isotopes?
-atoms of the same element with the same number of protons, but different numbers of neutrons and hence different mass numbers
what is meant by mass number?
-the total number of protons and neutrons in the nucleus of an atom
-different isotopes of the same element have different mass numbers
what is meant by relative atomic mass ( Ar )?
-the mean mass of the atoms of an element compared with 1/12 of the mass of a carbon-12 atom
-it is an average of the mass numbers of all the different isotopes of that element
how is relative atomic mass calculated?
-Ar = ( isotopic mass x % abundance ) + ( isotopic mass x % abundance ) / 100
-multiply the mass number of each isotope by its relative abundance
-add them all together
-divide by 100 if abundance is in %
why is the relative atomic mass not always a whole number?
-different isotopes of the same element have different mass numbers and the relative atomic mass is an average of the mass numbers of all these isotopes
what was Bohr’s theory?
-an atom has a small, positively charged central nucleus, orbited by electrons at fixed energy levels ( i.e. distances from the nucleus ) which are split into subshells
what is the structure of an atom?
-a small, positively charged central nucleus which contains protons and neutrons, orbited by electrons in shells which are made up of subshells ( s, p, d )
what is the nucleus of an atom?
-the positively charged central core of an atom, containing protons and neutrons and nearly all of its mass
what are the relative masses and charges of protons, neutrons and electrons?
-proton = 1, +1
-neutron = 1, 0
-electron = 0.0005 or 1/1836, -1
how can the number of protons, neutrons and electrons of an element be calculated?
-protons = atomic number
-neutrons = difference between mass ( big ) number and atomic ( small ) number
-electrons = number of protons
how are elements arranged in the periodic table?
-in order of increasing atomic number i.e. proton number, in rows called periods ( period number = number of electron shells )
-elements with similar properties are placed in the same vertical columns called groups ( group number = number of outer shell electrons )
-the order of elements is due to their atomic number, but an element’s position is due to its electronic structure and is dependent on its outermost electron shell
-metals on the left, nonmetals on the right and transition metals in the middle
-split into three blocks = s-block ( groups 1 and 2, along with H and He ), p-block ( groups 3-8 except He ) and d-block ( transition metals )
-s, p and d blocks indicate the subshell being filled with electrons as you go across the table, and the position of an element in a block is determined by the highest subshell occupied by electrons in that element
what is an electron shell?
-a group of atomic orbitals with the same principal quantum number, n
-also known as a main energy level
what is a subshell?
-a group of orbitals of the same type within a shell ( s, p or d )
what is an orbital?
-a region within an atom that can hold up to 2 electrons with opposite spins ( s, p or d )
-an s orbital is spherical in shape whereas a p orbital is dumb-bell shaped
how many electrons can each subshell hold and why?
-s subshell = 2, because it has 1 orbital
-p subshell = 6, because it has 3 orbitals
-d subshell = 10, because it has 5 orbitals
-f subshell = 14, because it has 7 orbitals
how many electrons can the first four shells hold and why?
-1st shell = 2, because it has an s subshell ( 1s )
-2nd shell = 8, because it has s and p subshells ( 2s, 2p )
-3rd shell = 18, because it has s, p and d subshells ( 3s, 3p, 3d )
-4th shell = 32, because it has s, p, d and f subshells ( 4s, 4p, 4d, 4f )
what is the order of filling orbitals?
-1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^10, 4p^6 etc.
what are the rules for arranging electrons in subshells?
-start at the lowest shell and add electrons one at a time
-fill each subshell before starting on the next ( the Aufbau principle )
-fill the the 4s subshell before the 3d subshell because 3d is higher in energy when occupied
-fill each orbital singly in a subshell before pairing electrons
-paired electrons have opposite spins to reduce repulsion, so they are shown as arrows pointing in opposite directions
what is a compound?
-a substance that is composed of two or more separate elements
-e.g. H2O, NaCl, CO2
what is a molecule?
-a group of atoms chemically bonded together, representing the smallest fundamental unit of a chemical compound that can take part in a chemical reaction
-e.g. H2O ( compound ), Cl2 ( element ), C2H6O ( compound )
what is an ion?
-an atom or molecule with a net electric charge due to the gain or loss of one or more electrons
-e.g. Cl^-, CO3^2-, Na^+
what is the formula of a nitrate ion?
-NO3^-
what is the formula of a carbonate ion?
-CO3^2-
what is the formula of a sulfate ion?
-SO4^2-
what is the formula of a hydroxide ion?
-OH^-
what is the formula of an ammonium ion?
-NH4^+
what is a monoatomic ion?
-formed by the gain or loss of electrons to the valence ( outermost ) shell in a single atom
-e.g. chloride, sodium, calcium
how is the charge on a monoatomic ion linked to its position in the periodic table?
-group 1 = 1+
-group 2 = 2+
-group 3 = 3+
-group 5 = 3-
-group 6 = 2-
-group 7 = 1-
why can’t the charges on ions of transition metals be predicted?
-they can form several stable ions
-e.g. iron can be Fe^2+ or Fe^3+
what is a molecular ion?
-formed by the gain or loss of elemental ions such as a proton, H+
-e.g. nitrate, carbonate, sulfate, hydroxide, ammonium
what is meant by relative formula mass ( Mr )?
-the sum of the relative atomic masses of all the atoms in a formula
what is meant by mole ( n )?
-the amount of any substance containing 6.02 x 10^23 particles
what is the mass of one mole of an element equal to?
-its relative atomic mass in g
-e.g. 24.3 g of magnesium contains 1 mole of atoms
what is the formula linking mass, moles and relative formula mass?
-m = n x Mr
-mass ( m ) in g = moles ( n ) in mol x relative formula mass ( Mr )
what is the Avogadro constant?
-the number of atoms, molecules or ions in one mole of a substance
-the value of the constant is 6.02 x 10^23
what is the formula linking the number of particles, the Avogadro constant and moles?
-P = n x Av
-number of particles ( P ) in atoms, molecules or ions = moles ( n ) in mol x the Avogadro constant ( Av )
what is meant by molar mass?
-the mass per mol of a substance in g mol^-1
-it is equal to its relative formula mass
what is the formula linking mass, moles and molar mass?
-m = n x M
-mass ( m ) in g = moles ( n ) in mol x molar mass ( M ) in g mol^-1
what is meant by empirical formula?
-the simplest / smallest whole number ratio of the atoms of each element in a compound
how is empirical formula calculated?
-calculate moles of each element by mass ( if in %, do mass in 100g ) / Ar
-divide moles by the smallest value to get a ratio
-adjust to make the ratios whole numbers by rounding up or doubling to get rid of halves
what is meant by molecular formula?
-the actual number of atoms of each element in a compound
how is molecular formula calculated?
-calculate Mr of empirical formula
-divide relative molecular mass by Mr of empirical formula
-multiply this by empirical formula
what is meant by stoichiometry?
-the ratio, of the amount in moles, of each substance in a chemical reaction ( the balancing numbers in front of each substance )
how is the mass ( in g ) of a substance calculated from the mass of another substance in the same reaction?
-balance equation to find molar ratio
-calculate moles of given by mass / Ar
-calculate moles of asked using molar ratio
-calculate Mr of asked
-calculate mass of asked by moles x Mr
what is meant by concentration?
-a measure of the amount of solute dissolved per unit of solvent
how is concentration calculated in mol dm^-3?
-c = n / V
-concentration ( c ) in mol dm^-3 = amount ( n ) in mol / volume ( V ) in dm^3
how is concentration calculated in g dm^-3?
-c = m / V
-concentration ( c ) in g dm^-3 = mass ( m ) in g / volume ( V ) in dm^3
how is cm^3 converted to dm^3?
-/ 1000
what is meant by percentage yield?
-the actual amount of yield worked out as a percentage of the theoretical yield
how is percentage yield calculated?
-actual amount ( moles or mass ) / theoretical amount ( moles or mass ) x 100
what is meant by theoretical yield?
-the expected amount of product from a reaction calculated from the balanced equation
why is the actual yield usually lower than the theoretical yield?
-the reaction may be incomplete ( hasn’t finished / reaches equilibrium )
-other side reactions may have occurred
-purification of the product may have resulted in loss of product
what is meant by first ionisation energy?
-the energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions
what is the equation for the first ionisation energy of lithium?
-Li(g) → Li+(g) + e-
what is meant by second ionisation energy?
-the energy required to remove the outermost electron from a 1+ ion
what is the equation for the second ionisation energy of lithium?
-Li+(g) → Li^2+(g) + e-
what factors affect ionisation energy?
-atomic number = the more protons in the nucleus, the greater the nuclear charge
-therefore, attraction of electrons to the nucleus increases and atomic radius decreases, so electrons are closer to the nucleus and require more energy to remove
-atomic size = the greater the atomic size, the greater the distance between the nucleus and the outer electrons
-therefore, attraction of electrons to the nucleus decreases, so less energy is required to remove them
-shielding = the more electrons between the nucleus and the outer electrons, the lower the effective nuclear charge because the attraction of the outer electrons to the nucleus is shielded ( reduced ) by the inner shell electrons
-therefore, atomic radius increases, so electrons are further away from the nucleus and require less energy to remove
what is the general trend in first ionisation energy down a group and why?
-first ionisation energy decreases down a group
-this is because there are more inner shell electrons as you go down a group, so the attraction of the outer electrons to the nucleus is shielded by the electrons in between
-therefore, atomic radius increases, so electrons are further away from the nucleus and require less energy to remove
what is the general trend in first ionisation energy across a period and why?
-first ionisation energy increases across a period
-this is because the nuclear charge increases as each successive element has one more proton than the last
-therefore, attraction of electrons to the nucleus increases and atomic radius decreases, so electrons are closer to the nucleus and require more energy to remove
what other factors affect ionisation energy?
-the amount of energy required to remove successive electrons increases because the overall positive charge has increased as well as the nuclear attraction, so atomic radius decreases
-there is a large increase in ionisation energy when an electron is removed from the shell below because these electrons are closer to the nucleus and there is less shielding
how do successive ionisation energies show which group an element is in?
-a large difference between two successive ionisation energies ( e.g. 2nd and 3rd ) shows the element is in the group number before the big jump ( e.g. group 2 )
-this is because the outer electrons are relatively easy to remove, while one from an inner shell which is closer to the nucleus is much more difficult to remove due to a stronger attraction and less shielding
what is meant by electron affinity?
-the energy change when an electron is added to each atom in 1 mole of gaseous atoms
what is the equation for the first electron affinity of fluorine?
-F(g) + e- → F-(g)
what is the general trend in first electron affinity down a group and why?
-first electron affinity decreases down a group
-this is because the shell to which the electron is being added is further from the nucleus and has more shielding from inner shells
-however, F and O are exceptions to this trend
what is meant by electronegativity?
-the ability of an atom to attract the bonding electron pair in a covalent bond
what is the general trend in electronegativity up a group and why?
-electronegativity increases up a group
-this is because the bonding electron pair in the covalent bond will be closer to the nucleus, attracting it
what is the general trend in electronegativity across a period and why?
-electronegativity increases across a period
-this is because the nuclear charge attracting the bonding electron pair is greater, but the number of electron shells shielding the bonding pair from the nucleus is the same
what is the general trend in atomic radius down a group and why?
-atomic radius increases down a group
-this is because the number of electron shells and inner shell electrons increases, so the shielding effect increases which reduces the attraction between the nucleus and the outer electrons, allowing the atomic radius to increase
what is the general trend in atomic radius across a period and why?
-atomic radius decreases across a period
-this is because the nuclear charge increases but the number of electron shells remains the same, so the outer electrons are more strongly attracted to the nucleus
what are the periodic trends in ionic radii?
-ionic radii increases down a group
-ionic radii decreases across a period
-cations are smaller than the atom they formed from
-anions are larger than the atom they formed from
-in the same period, the anions that form from the nonmetals are larger than the cations that form from the metals, as they have an extra electron shell
what is an ionic bond?
-the strong electrostatic attraction between oppositely charged ions
what is the structure of an ionic compound?
-between a metal and a nonmetal
-the metal loses electrons to become a positive cation, and the nonmetal gains electrons to become a negative anion
-giant, regular 3d lattice of alternating positive and negative ions
-the electrostatic forces between the ions act in all directions and keep the structure together
what factors affect the strength of an ionic bond?
-ionic charge = the greater the charge on the ions, the stronger the bond
-ionic radius = the smaller the size of the ions, the stronger the bond
what is a covalent bond?
-the electrostatic attraction between a shared pair of electrons and the nuclei of the bonded atoms
what is the structure of a covalent compound?
-between nonmetals
-atoms share a pair of electrons
-the electron pair is attracted to both nuclei and is localised between them
-generally, each atom contributes 1 electron to the pair, but a covalent bond consisting of an electron pair derived from one of the atoms is called a dative ( coordinate ) covalent bond
-atoms with lone ( non-bonded ) pairs of electrons are able to form dative covalent bonds with other atoms which are electron deficient
what factors affect the strength of a covalent bond?
-bond length = the shorter the bond length, the stronger the bond
-number of electron pairs / bonds = the more electron pairs in the bond, the stronger the bond
what makes a covalent compound tetrahedral?
-when there are four pairs of outer electrons around a central atom ( four bonds ), the orbitals containing the pairs repel as far away as possible, forming bond angles of 109.5°
-e.g. methane which consists of 1 carbon atom bonded to 4 hydrogen atoms
what is a metallic bond?
-the strong electrostatic attraction between positive metal ions and delocalised electrons
what is the structure of a metallic compound?
-between metals
-giant, regular 3d lattice of positive ions surrounded by a sea of delocalised electrons ( the electrons from the outer shell of the metal atoms that are free to move throughout the structure )
-electrostatic attraction between nuclei of metal cations and delocalised electrons
what are the properties of metals and why?
-electrical and thermal conductivity due to the delocalised electrons which are free to move
-high melting and boiling points due to strong electrostatic attractions between positive ions and electrons
-malleability = can be shaped as regular structure means layers of positive ions slide over each other and the delocalised electrons move with the layers, so strong metallic bonds remain intact
-ductility = can be pulled into wires as layers of positive ions slide over each other and the delocalised electrons move with the positive ions, so strong metallic bonds remain intact
why does the metallic character of elements increase down a group?
-the atoms increase in radius and have more shielding electrons, so the electrostatic attraction between the nucleus and the outer electrons is weaker
-therefore, the electrons can move more freely so the elements can conduct heat and electricity more effectively, and the outer electrons can be lost more readily to form cations in chemical reactions
what are intermolecular forces?
-the weak forces between molecules caused by either permanent or induced dipoles that hold together simple covalent molecules
what are the three types of intermolecular forces and when do they occur?
-London dispersion ( van der Waals / temporary dipole-induced dipole ) forces = between all atoms
-dipole-dipole interactions = between polar molecules ( which have a 𝛿− and 𝛿+ side )
-hydrogen bonds = between a hydrogen atom covalently bonded to oxygen, nitrogen or fluorine and another very electronegative atom
what is the order of intermolecular forces from weakest to strongest?
-London forces
-dipole-dipole interactions
-hydrogen bonds
how do London forces form?
-electrons are moving randomly within the shells of a molecule or atom
-this can cause an uneven distribution in the molecule, causing partial charges ( 𝛿− and 𝛿+ ) to develop and an instantaneous, temporary dipole to form, which can induce a temporary dipole in a nearby molecule
-this results in a weak attraction ( a London force )
how does the number of electrons affect London forces?
-the more electrons a molecule has, the more likely this process occurs, so the stronger the London forces
-when two molecules have a similar number of electrons, the one with the shape that has greatest surface area will have greater London forces
how do dipole-dipole interactions form?
-polar molecules ( e.g. HCl ) have permanent dipoles due to the much greater electronegativity of one atom ( e.g. the chlorine atom ) and the fact that the molecule is not symmetrical
-e.g. the chlorine atom has a partial negative charge ( 𝛿− ) and the hydrogen atom has a partial positive charge ( 𝛿+ )
-therefore, the oppositely charged ends of two polar molecules are attracted to each other, forming a weak permanent dipole-dipole interaction between the molecules
how does the difference in electronegativity affect dipole-dipole interactions?
-the greater the difference in electronegativity between the two atoms, the stronger the permanent dipole-dipole interaction
why are oxygen, nitrogen and fluorine the only atoms that can form hydrogen bonds?
-they are small and highly electronegative, which means they pull pairs of electrons towards them
how do hydrogen bonds form?
-the large difference in electronegativity when hydrogen is bonded to one of three very electronegative elements ( i.e. oxygen, nitrogen or fluorine ) results in a polar bond, with significant charge separation and a dipole moment
-the electron deficient hydrogen on one molecule then forms a bond with a lone pair of electrons on a neighbouring molecule ( O, N or F )
what is the difference between ice and liquid water?
-ice is less dense than water because when water freezes, its molecules form a giant lattice structure in which they are further apart from each other than in their liquid state, therefore less mass per unit volume
what is the trend in melting and boiling points down groups 7 and 0 and why?
-as the molecules ( group 7 ) or atoms ( group 0 ) become larger as you go down both groups, the number of electron shells increases
-although the particles are non-polar and do not have a permanent uneven distribution of electrons, they can have a temporary uneven distribution of electrons
-this means as you go down the groups, with the increase in electrons and distance over which they can move, the bigger the possible temporary dipoles and therefore the bigger the dispersion forces
-the resulting increased London forces require more energy to overcome during melting or boiling, so melting and boiling points increase down groups 7 and 0
what is the general trend in melting and boiling points across a period and why?
-melting and boiling points increase across groups 1 to 4 and decrease across groups 5 to 0 as the type of bonding changes from metallic to covalent across the period
-the melting and boiling points of metals increase as the metallic bond strength increases because the positive metal ions get smaller across the period, their charge increases and there are more delocalised electrons, so there is a stronger attraction between the cations and the delocalised electrons
-the elements from group 4 ( carbon and silicon ) both have very high melting and boiling points as they have giant covalent structures with many strong covalent bonds
-the elements from groups 5 to 0 are simple molecules / atoms with only weak London forces between them
what are the four types of structures and their properties?
-giant ionic lattice = high melting and boiling points, conduct heat and electricity when dissolved in solution or molten, soluble in water
-simple molecular covalent = low melting and boiling points, don’t conduct heat or electricity
-giant covalent lattice = very high melting and boiling points, conduct heat and electricity due to delocalised electrons ( e.g. in graphite )
-giant metallic lattice = quite high melting and boiling points, conduct heat and electricity due to delocalised electrons
what is meant by oxidation number / state?
-a measure of the number of electrons that an atom uses to bond with an atom of another element
what is the oxidation number of a neutral element?
0
-e.g. H2, F2, Na, O2, C ( diamond )
what is the oxidation number of a monoatomic ion?
-equal to the charge on the ion
-e.g. Na^+ = +1, Cl^- = -1, S^2- = -2
what is the sum of all oxidation numbers in a neutral compound?
0
-e.g. sum in HCl = 0
what is the sum of all oxidation numbers in a polyatomic ion?
-equal to the charge on the ion
-e.g. sum in CO3^2- = -2
what is the oxidation number of a group 1 element in a compound?
+1
-e.g. Na in NaCl = +1
what is the oxidation number of a group 2 element in a compound?
+2
-e.g. Ba in BaBr2 = +2
what is the oxidation number of a group 3 element in a compound?
+3
-e.g. Al in AlCl3 = +3
what is the oxidation number of hydrogen in most compounds?
+1
-e.g. H in H2SO4 = +1
what is the oxidation number of oxygen in most compounds?
-2
-e.g. O in MgO = -2
what is the oxidation number of fluorine in every compound?
-1
-e.g. F in NaF = -1
what is the oxidation number of chlorine, bromine and iodine in compounds without oxygen?
-1
-e.g. Cl in HCl = -1
what is the oxidation number of transition metals in compounds?
-can vary, so is usually mentioned in the compound’s name as roman numerals
-e.g. Fe in Fe2O3 = +3 but Fe in FeO = +2
what does a redox reaction involve?
-oxidation and reduction
-e.g. displacement reactions
what happens when an element is oxidised?
-it loses electrons
-its oxidation number increases
what happens when an element is reduced?
-it gains electrons
-its oxidation number decreases
what happens in a metal displacement reaction?
-a more reactive metal displaces a less reactive metal from a metal salt
-the more reactive metal atoms lose electrons to form ions, so are oxidised
-each metal ion gains electrons to form atoms, so is reduced
what happens in a halogen displacement reaction?
-a more reactive halogen ( further up group 7 ) displaces a less reactive halogen from a halide salt
-the more reactive halogens gain one electron per atom to form halide ions, so are reduced
-each halide ion loses an electron to form halogens, so is oxidised
what happens when period 2 and 3 elements react with oxygen?
-they form oxides, often by being burnt in pure oxygen
-e.g. Li2O, Na2O
when do the oxides tend to form alkaline solutions and why?
-if the metal oxides dissolve in water as they form metal hydroxides
-e.g. sodium oxide reacts to form sodium hydroxide
when do the oxides tend to form acidic solutions and why?
-if the nonmetal oxides dissolve in water as they form acids
-e.g. sulfur dioxide reacts to form sulfurous acid
what are the reactions of metals with water?
-some metals ( e.g. those in groups 1 and 2 ) react directly with water to form hydroxides
-the metals in group 1 react immediately on contact with water, fizzing as hydrogen is produced, dissolving to form an alkaline solution, and sometimes catching on fire or exploding
-the metals in group 2 react in a similar but slower way, e.g. magnesium needs to be reacted with steam to form hydrogen
-less reactive metals don’t tend to react directly with water, but may react with water and oxygen to corrode
what are the reactions of metals with dilute acids?
-metals that are more reactive than hydrogen ( potassium, sodium, lithium, calcium, magnesium, aluminium, zinc and iron ) will react with dilute acids to form a salt and hydrogen gas
-the salt forms when a metal ionises and metal ions replace hydrogen ions in the acid
-the salt formed depends on the dilute acid used ( sulfuric acid = sulfate, hydrochloric acid = chloride )
what are the trends in metal reactivity and why?
-as you go down a group, the metals become more reactive with oxygen, water and dilute acids because the atoms become larger in size
-as you go across a period, the metals become less reactive with oxygen, water and dilute acids because more electrons have to be lost to form metal ions
what is a transition metal?
-a d-block element which forms at least one stable ion with an incomplete d-subshell
what are some examples of uses of nonmetal oxides?
-boron trioxide ( B2O3 ) = glass manufacture, e.g. optical fibres
-carbon dioxide ( CO2 ) = to make drinks fizzy and in fire extinguishers to prevent oxygen gas reaching flames
-nitrous oxide ( N2O ) = pain relief, e.g. during childbirth
-silicon dioxide ( SiO2 ) = food additive to stop powders sticking together
-phosphorus pentoxide ( P2O5 ) = removes water from organic molecules in the chemical industry
-sulfur dioxide ( SO2 ) = manufacture of sulfuric acid ( contact process )
what are some examples of uses of metal oxides and hydroxides?
-aluminium oxide ( Al2O3 ) = in abrasive paper as its giant ionic lattice structure makes it very hard
-magnesium hydroxide ( Mg(OH)2 ) = antacid medicine as it neutralises excess HCl in the stomach
-sodium hydroxide ( NaOH ) = drain cleaner as it will react and breakdown fats and oils from food waste
what are some examples of uses of metal salts?
-sodium chloride ( NaCl ) = food industry as flavouring and a preservative, and raw material in chemical industry to make hydrogen, chlorine and sodium hydroxide through electrolysis as it’s readily avaliable and very soluble in water
-potassium sulfate ( K2SO4 ) = fertiliser as it’s very soluble and contains minerals needed by plants
-magnesium sulfate ( MgSO4 ) and sodium sulfate ( Na2SO4 ) = a drying agent as their giant ionic crystalline structure can absorb water molecules
-calcium sulfate ( CaSO4 ) = key component of plaster used to cover walls in buildings
what are some examples of uses of transition metals?
-copper = in electrical wiring as its metallic structure allows electrons to move freely and so conduct electricity effectively
-vanadium = added to steel as vanadium steels are very hard and resistant to wearing so can be used in engines
-titanium = in aircraft manufacture as it’s as strong as steel but much less dense
