Nuclear Model of the Atom PDF

Nuclear model of the atom Electrons (e-) Atoms are made up of 3 types of particles: 1. Neutrons: have no charge but have mass. Found in the nucleus. 2. Protons: have positive charge and same mass as neutrons. Found in the nucleus. 3. Electrons: have negative charge and a mass 1/1836 that of protons. Found in all the space surrounding the nucleus. Nucleus For a neutral atom No of protons = No of electrons 1 atomic mass unit= 1.67 x 10-27 kg Trinity College Dublin, The University of Dublin 1. Atomic number (Z) Z = number of protons in nucleus For a Neutral atom: Also the number of electrons orbiting the nucleus 2. Mass number (A) A = number of protons + neutrons in the nucleus No of neutrons = A – Z Atomic Properties Let’s consider an example: 1. Z =noofprotons=6 2. No of protons = no of electrons = 6 3. Noofneutrons=A–Z=12–6=6 Mass number Atomic number No of protons = No of electrons = No of neutrons = Trinity College Dublin, The University of Dublin Atomic number Element symbol Element name Mass number 1. No of protons = 2. No of electrons = 3. No of neutrons = Trinity College Dublin, The University of Dublin Atomic number Element symbol Element name Mass number One more example: Z = no of protons = 29 = no of electrons A = 63.55 ≈ 64 No of neutrons = A – Z = 64 – 29 = 35 1. No of protons = 2. No of electrons = 29 3. No of neutrons = 35 29 Trinity College Dublin, The University of Dublin Things we now know about atoms An atom contains a very small nucleus composed of protons and neutrons. Protons and neutrons are much heavier than electrons; most of the mass is in the nucleus. Electrons occupy almost all of an atom’s volume. Atoms of one element differ in properties from atoms of all other elements. An element consists of only one type of atom. A macroscopic sample of an element contains an incredibly large number of atoms, all of which have identical chemical properties. A compound consists of atoms of two or more elements in a whole-number ratio. If the nucleus was the size of a blueberry, the atom would be a football stadium. Trinity College Dublin, The University of Dublin Periodic table Trinity College Dublin, The University of Dublin Birth of periodic table Antoine Lavoisier (1789) Grouped elements into gases, metals, non-metals and earth elements. Wolfgang Dobereiner (1829) grouped atoms into “triads”, groups of similarly behaving atoms For example: Lithium, sodium and potassium Dmitri Mendeleev (1869) Listed elements in order of atomic weights Trinity College Dublin, The University of Dublin Birth of periodic table-Mendeleev Saw trends in physical and chemical properties based on atomic weight (AR). Placed elements into rows of increasing mass and columns of similar reactivity. Gaps appeared, so he predicted the mass and properties of these ‘undiscovered’ elements Soon afterwards they were actually discovered Trinity College Dublin, The University of Dublin Modern periodic table Now arranged in order of increasing atomic number (Z) Vertical columns are called groups Ø Elements in a group exhibit trends in chemical reactivity. Horizontal rows are called periods. Ø A number of trends are found within periods. There are 4 blocks, based on where the element’s outer electrons lie. Ø s block, p block, d block, f block. Looking at an element’s position can tell you about its chemical and physical properties. Trinity College Dublin, The University of Dublin Groups Trinity College Dublin, The University of Dublin Periods Transition metals Trinity College Dublin, The University of Dublin Groups of the periodic table Group 1: Alkali metals Ø Soft silvery metals, not found in elemental state in nature. Ø Melt at low temperature. Ø Produce H2 on reaction with water. Ø Reactivity increases down the group. Ø Each have a single electron in their outermost shells Trinity College Dublin, The University of Dublin Groups of the periodic table Group 2: Alkaline earth metals Ø Soft metals. Ø Often found in nature as oxides (Calcium oxide, barium oxide) Ø Commonly found in nature (magnesium and calcium ) Ø Used in radiation therapy (Radium) Ø Each have two electrons in their outermost shells Trinity College Dublin, The University of Dublin Groups of the periodic table Group 17: Halogens Ø Very reactive Ø Fluorine and chlorine are gases Ø Bromine is liquid and iodine a solid Ø Fluoride used in toothpaste, chloride use for disinfection, iodide has several important functions in the body (thyroid gland) Trinity College Dublin, The University of Dublin Groups of the periodic table Group 18: Noble gases Ø Colourless, odourless gases. Ø Full outer-shell (closed shell) electron-quite unreactive Ø Low temperature cryogenics (helium) Ø Absorb and emit electromagnetic radiation (neon lamps) Trinity College Dublin, The University of Dublin d block elements Transition Elements Ø Bridge the active groups 1 – 2 and 12 – 18. Ø Can form compounds in several oxidation states Ø Form highly coloured compounds Ø Fe, Co, Ni can display ferromagnetism Trinity College Dublin, The University of Dublin f block elements Lanthanides and Actinides (inner transition elements) Ø Heavy metals Ø Can form compounds in several oxidation states Ø Actinides are radioactive in nature (Nuclear fuel, imaging) Trinity College Dublin, The University of Dublin Trends in the periodic table 1. Atomic radii 2. Ionic radii 3. Electronegativity 4. Ionisation energy Trinity College Dublin, The University of Dublin 1. Trends in atomic radii Atomic radius is the size of an atom of a given element. increased nuclear charge-> outermost electrons are pulled closer to the nucleus. Trinity College Dublin, The University of Dublin 1. Trends in atomic radii more electron shells -> outermost electrons get further from nucleus. Atomic radius decreases across the table and increases down the periodic table. Trinity College Dublin, The University of Dublin Atomic radius increases 2. Trends in ionic radius Net positive/negative charge due to loss/gain of an electron Ionic radius is the size of an ion for a given element Metalsbecomesmalleruponionisation(theyloseelectrons) Ca > Ca2+ Non-metals(right)getbigger(theygainelectrons.) O < O2- Decreases across the table and increases down the periodic table. Trinity College Dublin, The University of Dublin 3. Trends in electronegativity Electronegativity : ability of an atom to attract electrons. Left of table: valence shells less than half full -> these atoms (metals) tend to lose electrons (low electronegativity). Right of table: valence shells more than half full -> these atoms (nonmetals) tend to gain electrons (high electronegativity). Down table: number of energy levels (and distance from nucleus to outer orbitals) increase. Shielding weakens nuclear attraction (ability of atom to attract electrons) Increases across the periodic table and decreases down the table. Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin 4. Trends in ionisation energy Ionisation energy: energy needed to remove one electron from the outermost shell of an atom. Electronegative elements and noble gases: don’t want to lose an electron->high ionisation energies. Elements with many full inner shells: interact poorly with outermost electrons->low ionisation energies. Increases across the periodic table and decreases down the table. Atomic structure Trinity College Dublin, The University of Dublin Atomic structure and periodicity Atomic Number gives the number of protons and electrons in an element. Octet Rule: During a chemical reaction, an element will try to attain the same number of electrons as the nearest noble gas. Atoms try to gain, lose or share electrons in order to have the same number of electrons as a noble gas. The distribution of electrons within the levels and sublevels—control these chemical and physical properties. This is electron configuration Trinity College Dublin, The University of Dublin Quantum numbers We can describe each electron individually through 4 quantum numbers. Pauli Exclusion Principle: These are unique and cannot be shared. Electrons within an atom are confined to a set of energy levels. The ‘address’ of an electron within the atom is described using the following 4 quantum numbers: n Principal Energy Level l Subshell ml Orbital ms Spin Trinity College Dublin, The University of Dublin 1. Principal energy levels (n) Each Principal Energy Level (or electron shell) has a fixed maximum number of electrons. The higher the energy level, the more electrons it can hold: n = 1 n = 2 n = 3 etc... Holds 2 electrons Holds 8 electrons Holds 18 electrons n=3 n=2 n=1 Trinity College Dublin, The University of Dublin 2. Energy subshells (l): secondary quantum number Ø Also known as Orbital, angular or azimuthal quantum number ØShells are subdivided into subshells. The number of subshells within each shell is the principal energy level (n) of that shell: n = 1 n = 2 n = 3 etc... has 1 subshell (s) has 2 subshells (s, p) has 3 subshells (s, p, d) Trinity College Dublin, The University of Dublin 3. Atomic orbitals (ml) The subshells are divided into atomic orbitals. Each atomic orbital can hold 2 electrons. Each electron in the pair must have opposite spin (spin can be ‘up’ or ‘down’). s subshells have 1 orbital (2 electrons) p subshells have 3 orbitals (6 electrons) d subshells have 5 orbitals (10 electrons) f subshells have 7 orbitals (14 electrons) etc.... Trinity College Dublin, The University of Dublin 4. Spin quantum number (ms) Each electron in a pair must have opposite spin (spin can be ‘up’ or ‘down’). Clockwise rotation around the axis of the electron m s = + !" This is represented by ↑ Anti-clockwise rotation around the axis of the electron m s = - !" This is represented by ↓ ↑↓ Trinity College Dublin, The University of Dublin n=4 n=1 Ø n=1 (1st shell) Ø 1 subshell = s Ø n=4 (4th shell) Ø 4 subshells = s, p, d and f Ø Can hold 32 electrons n=2 n=3 Ø n=2 (2nd shell) Ø 2 subshells = s and p Ø Can hold 8 electrons à 2+6 Ø n=3 (3rd shell) Ø3 subshells = s, p and d Ø Can hold 18 electrons à 2+6+10 Trinity College Dublin, The University of Dublin Electron structure and periodicity The principal energy level (n) corresponds to the row of the periodic table n=1 n=2 n=3 n=4 n=5 n=6 n=7 5 pairs/subshell 1 pairs/subshell 3 pairs/subshell The electrons within subshells pair up into atomic orbitals (ml) s block 2 electrons d block 10 electrons p block 6 electrons Each of the two electrons in an orbital have opposite spins (ms) Subshells (l) correspond to blocks of the periodic table. Trinity College Dublin, The University of Dublin Shapes of atomic orbitals ↑↓ s s orbitals Ø Shaped like a sphere Ø There is one s orbital within the s subshell, it can hold 2 electrons p orbitals Ø Have two lobes Ø Shaped like a dumbbell Ø In the p subshell, there are 3 p orbitals of equal energy ↑↓ ↑↓ ↑↓ 𝑝# 𝑝$ 𝑝% Trinity College Dublin, The University of Dublin Shapes of atomic orbitals d orbitals (and higher) Ø Can have complicated shapes ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 𝑑#$ 𝑑#% 𝑑$% 𝑑#! &$! 𝑑%! Ø Electron orbitals with same energy orbitals is known as degenerate orbitals Ø p has 3 degenerate orbitals, d has 5 Trinity College Dublin, The University of Dublin Electron pairing rules Hund’s Rule: Whenever 2 or more orbitals of equal energy are available (such as px, py, pz), the electrons will occupy them singly before pairing up. ↑↓↑ ↑↑↑ 𝑝# 𝑝$ 𝑝% 𝑝# 𝑝$ 𝑝% A p subshell with 3 electrons will have one electron in each p orbital. This is explained in terms of energy. There is a pairing energy for pairing +1/2 and -1/2 electrons. This is a penalty when there are free orbitals of the same energy. Wrong! Correct J Trinity College Dublin, The University of Dublin Electron pairing rules Pauli’s Exclusion Principle: Each orbital can hold one pair of electrons; but the electrons will have opposite spins. ↑↑ ↑ ↑ ↑↓ ↑ ↑ 𝑝# 𝑝$ 𝑝% 𝑝# 𝑝$ 𝑝% A p subshell with 4 electrons will have 3 spin-up electrons and 1 spin- down. Wrong! Correct J This is a quantum explanation. An electron is described by the quantum numbers (n, l, ml , ms). They are unique to one electron and cannot be duplicated. Trinity College Dublin, The University of Dublin The Aufbau principle Electrons occupy the lowest available energy level, in order to be as close to the nucleus as possible. The order in which these orbitals are filled are shown below: 1s 2s 2p 3s 3p 4s 4p 4d 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s Note that the 4s orbital is filled before the 3d orbital. 3d 4f 5s 5p 6s 6p 6d 7s 7p 5d 5f Trinity College Dublin, The University of Dublin Writing electron configurations Electronic configuration: The arrangement of electrons within an atom. Describes how many electrons are present in each orbital of that atom. The electronic configuration of Hydrogen is given below: No of electrons within that subshell Hydrogen: 1 𝑠% Subshell energy level Principal Trinity College Dublin, The University of Dublin The n = 1 energy level The n=1 energy level has only a single s orbital, the 1s orbital. Hydrogen and Helium are both n = 1 elements. Their electronic configurations are: Hydrogen: 1 𝑠% ↑ n=1 Helium: 1 𝑠& ↑↓ 1s 2s 2p 3s 3p 4s 5s 5p 6s 7s 7p 3d 4p 4d 4f 5d 5f 6p 6d Trinity College Dublin, The University of Dublin The n = 2 energy level The n=2 energy level has both s and p subshells; called the 2s and the 2p. Carbon is an n=2 element, its electronic configuration is : Carbon:1𝑠&2𝑠&2p& 1 𝑠 & 2 𝑠 & 2 𝑝 '% 2 𝑝 (% ↑↓ ↑↓↑ ↑ 1s 2𝑠 2𝑝# 2𝑝$ 2𝑝% 1s 2s 2p 3s 3p 4s 4p 5s 5p 6s 6p 7s 7p 6 electrons 3d 4d 4f 5d 5f 6d n=2 Trinity College Dublin, The University of Dublin The n=3 energy level The n=3 energy level has s, p, and d subshells; called the 3s, the 3p and 3d subshells. Note: The 3d is only filled after the 4s orbital (3d filled when n=4)! Sulfur: 1s 2s 2p 1 𝑠& 2𝑠& 2p+ 3s& 3p, 1 𝑠& 2𝑠& 2𝑝'& 2 𝑝(& 2𝑝-& 3𝑠& 3𝑝'& 3𝑝(% 3𝑝-% n=3 16 electrons 3s 4s 5s 5p 6s 7s 7p 3p 3d 4p 4d 4f 5d 5f 6p 6d Trinity College Dublin, The University of Dublin Electron configuration examples Hydrogen atom (H): Z = 1 1𝑠 Helium atom (He): Z = 2 1𝑠 Lithium atom(Li): Z = 3 1𝑠 2𝑠 2𝑝 → it has one electron ↑ H: 1 𝑠% → it has two electrons ↑↓ He:1𝑠& → it has three electrons ↑↓ ↑ 2𝑝 2𝑝 2𝑝 #$% Li: 1 𝑠& 2𝑠% Li: 𝐻𝑒 2𝑠% 1s 2s 2p 3s 3p 4s 4p 5s 5p 6s 6p 7s 7p 3d 4d 4f 5d 5f 6d Trinity College Dublin, The University of Dublin Electron configuration examples Carbon atom (C): Z = 6 → it has six electrons 1𝑠 2𝑠 2𝑝 ↑↓ ↑↓ ↑ ↑ C:1𝑠&2𝑠&2𝑝'% 2𝑝(% & & & C:1𝑠 2𝑠 2𝑝 2𝑝# Nitrogen atom(Li): Z = 7 1𝑠 2𝑠 2𝑝 ↑↓ ↑↓ 2𝑝$ 2𝑝% → it has seven electrons ↑↑ 2𝑝# 2𝑝$ 2𝑝% N:1𝑠&2𝑠&2𝑝'% 2𝑝(%2𝑝-% ↑ N: 1 𝑠& 2𝑠&2𝑝. Trinity College Dublin, The University of Dublin Electron configuration examples Oxygen atom (O): Z = 8 → it has eight electrons 1𝑠 2𝑠 2𝑝 O:1𝑠&2𝑠&2𝑝'& 2𝑝(%2𝑝-% O: 1 𝑠& 2𝑠& 2𝑝, ↑↓ ↑↓ Neon atom(Ne): Z = 10 1𝑠 2𝑠 2𝑝 2𝑝$ 2𝑝% → it has ten electrons Ne: 1 𝑠& 2𝑠& 2𝑝'& 2𝑝(& 2𝑝-& ↑↓ ↑ ↑ 2𝑝# ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ Ne: 1 𝑠& 2𝑠& 2𝑝+ 2𝑝# 2𝑝$ 2𝑝% Trinity College Dublin, The University of Dublin Argon atom (Ar): Z = 18 1𝑠 2𝑠 2𝑝 → it has eighteen electron 1s 2s 2p 3s 3p 4s 4p 5s 5p 6s 6p 7s 7p ↑↓ ↑↓ ↑↓ 2𝑝# 3𝑠 3𝑝 ↑↓↑↓ ↑↓ ↑↓↑↓↑↓ 3d 4d 4f 5d 5f 6d Electron configuration examples 2𝑝$ 2𝑝% 3𝑝# 3𝑝$ 3𝑝% Ar: 1 𝑠& 2𝑠& 2p+ 3s& 3p+ Ar:[𝑁𝑒] 3s& 3p+ Trinity College Dublin, The University of Dublin The n=3 energy level Sulfur: 1 𝑠& 2𝑠& 2p+ 3s& 3p, 1 𝑠& 2𝑠& 2𝑝'& 2 𝑝(& 2𝑝-& 3𝑠& 3𝑝'& 3𝑝(% 3𝑝-% 1𝑠 2𝑠 2𝑝 3𝑠 3𝑝 ↑↓↑↓↑↓ ↑↓ ↑↓ ↑↓↑↓ ↑ ↑ 2𝑝# 2𝑝$ 2𝑝% 3𝑝# 3𝑝$ 3𝑝% 1s 2s 2p 3s 3p 4s 4p 5s 5p 6s 6p 7s 7p 3d 4d 4f 5d 5f 6d Trinity College Dublin, The University of Dublin Question Q. Write the electronic configuration of Sodium (Z=11). 3𝑠 ↑ 1𝑠 2𝑠 2𝑝 ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 2𝑝# 2𝑝$ 2𝑝% Na: 1𝑠& 2𝑠& 2p+ 3s% Na: 𝑁𝑒 3s% 1s 2s 2p 3s 3p 4s 4p 5s 5p 6s 6p 7s 7p 3d 4d 4f 5d 5f 6d Trinity College Dublin, The University of Dublin Electronic configurations of ions Ions are atoms which have lost or gained electrons. To write the configuration of an ion, we add or remove enough electrons from the atom to make up the ion’s charge. Eg.: Li+ has a charge of +1, so its configuration is equal to the configuration of lithium minus 1 electron. » Li = 1s2, 2s1 (Remove outermost electron) » Li+= 1s2= [He] A more complex example is O2-: » O = [He], 2s2, 2px2, 2py1, 2pz1 (add 2 e-) » O2-= [He], 2s2, 2px2, 2py2, 2pz2 = [Ne] Trinity College Dublin, The University of Dublin Quantum numbers We can describe each electron individually through 4 quantum numbers. Pauli Exclusion Principle: These are unique and cannot be shared. Electrons within an atom are confined to a set of energy levels. The ‘address’ of an electron within the atom is described using the following 4 quantum numbers: n Principal Energy Level l Subshell ml Orbital ms Spin Trinity College Dublin, The University of Dublin 1. Principal energy levels (n) Each Principal Energy Level (or electron shell) has a fixed maximum number of electrons. The higher the energy level, the more electrons it can hold: n = 1 n = 2 n = 3 etc... Holds 2 electrons Holds 8 electrons Holds 18 electrons n=3 n=2 n=1 Trinity College Dublin, The University of Dublin 2. Energy subshells (l): secondary quantum number Also known as Orbital, angular or azimuthal quantum number Shells are subdivided into subshells. The number of subshells within each shell is the principal energy level (n) of that shell: n = 1 n = 2 n = 3 etc... has 1 subshell (s) has 2 subshells (s, p) has 3 subshells (s, p, d) Trinity College Dublin, The University of Dublin 3. Atomic orbitals (ml) The subshells are divided into atomic orbitals. Each atomic orbital can hold 2 electrons. Each electron in the pair must have opposite spin (spin can be ‘up’ or ‘down’). s subshells have 1 orbital (2 electrons) p subshells have 3 orbitals (6 electrons) d subshells have 5 orbitals (10 electrons) f subshells have 7 orbitals (14 electrons) etc.... Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin 4. Spin quantum number (ms) Each electron in a pair must have opposite spin (spin can be ‘up’ or ‘down’). Clockwise rotation around the axis of the electron ms = + 1 22 ↑↓ Anti-clockwise rotation around the axis of the electron This is represented by ↑ This is represented by ↓ ms = - 1 Trinity College Dublin, The University of Dublin n=4 n=1 (1st shell) 1 subshell = s n=4 (4th shell) 4 subshells = s, p, d and f Can hold 32 electrons n=2 n=1 n=3 n=2 (2nd shell) 2 subshells = s and p Can hold 8 electrons → 2+6 n=3 (3rd shell) 3 subshells = s, p and d Can hold 18 electrons → 2+6+10 Trinity College Dublin, The University of Dublin Source: Silderberg, Chapter 8 Electron Configuration and Chemical Periodicity Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin The chemical bond A chemical bond is a connection between a pair of atoms formed by an exchange of their outermost electrons. To generate a chemical bond, electrons in the outermost shell may be shared between two atoms, or transferred from one atom to the other. If the new arrangement of electrons is more stable than that of the isolated atoms, a chemical bond is formed. Chemical reactions: loss, gain, or rearrangement of valence electrons (not core electrons) Trinity College Dublin, The University of Dublin Ionic vs covalent bonding An ionic bond is formed by the attraction of oppositely charged ions E.g. Na+ + Cl- → NaCl The bonding electrons are localised around the negative ion: “electron density” i.e. where the electrons are Na+ Cl- A covalent bond is formed by sharing electrons between two uncharged atoms e.g. H+H→ H2 The bonding electrons are shared between the two atoms: H H Trinity College Dublin, The University of Dublin Electronegativity and bonding Electronegativity = Ability of an atom to attract electrons. When two atoms with a big difference in electronegativity come together; they form an ionic bond. When two atoms with a small difference in electronegativity come together, they form a covalent bond. H + H → H2 Na+ + Cl- → NaCl Trinity College Dublin, The University of Dublin Ionic vs. covalent bonding In reality, chemical bonds are very rarely completely ionic or completely covalent! Instead, there is a spectrum of chemical bonding: No difference in electronegativity Covalent bonding Polar covalent bonding Big difference in electronegativity Ionic bonding Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin Drawing Lewis structures A Lewis Structure is a diagram of an element showing its valence electrons (outer shell electrons) drawn as dots. H Hydrogen Lewis structure We represent pairs of electrons in the same orbital as two dots, unpaired electrons are drawn as isolated, unpaired dots. Oxygen: 8 electrons Valence electrons ↑↓ ↑↓ ↑↓ ↑ ↑ O:1𝑠22𝑠22𝑝2 2𝑝12𝑝1 𝑥𝑦𝑧 An unpaired electron O A pair of electrons 2 pairs of valence electrons 2 unpaired valence electrons Trinity College Dublin, The University of Dublin Lewis structure – Ionic bonds 1. Lithium has 1 outer electron (wants to lose 1, valence = 1), while Bromine 7 outer electrons (wants to gain 1, valence =1): Li Br The formation of Lithium Bromide can be drawn using Lewis Structures (we draw a cross to highlight Lithium’s electron): Lix + Br Li+ xBr- 2. Calcium has 2 valence electrons while chlorine has 7 valence electrons; the formation of calcium chloride is therefore: Cl + xCax + Cl Cl - Ca2+ - xCl x Trinity College Dublin, The University of Dublin LiF Na2O Trinity College Dublin, The University of Dublin Lewis structures – covalent bonds Hydrogen has 1 valence electron but needs 2 to fill its outer shell. By sharing their electrons, two H atoms can fill their outer shells: H+xH HxH A pair of electrons between two atoms represents a bond. We can draw a line instead to represent this bonding electron pair: H+xH HH Fluorine has 7 outer electrons. By sharing their electrons, two F atoms can both have 8 electrons in their outer shell to form F2: FF F Fx F Fx xx xx xx xx +xx xx Trinity College Dublin, The University of Dublin Compounds Compounds can also be divided into molecular and ionic. » Molecular compounds have atoms bonded together into molecules Oxygen Hydrogen »Ionic compounds consist of ions (positively or negatively charged atoms in groups). Na+ Cl- O xygen Trinity College Dublin, The University of Dublin Ionic compounds-monatomic ions An ionic compound consists of 2 or more electrically charged atoms (ions) in which the overall charge is zero Eg: MgCl2→Mg2+ + 2Cl- Monoatomic cations keep the name of the element. If more than 1 cation is possible include the charge. Na+ = sodium Cu1+ = copper (I) Monoatomic anions change their ending to “-ide”. F- = fluoride Cl- = chloride S- = sulphide Cu2+ = copper (II) O2- = oxide Ionic nomenclature: Cations first, anion next CaS Calcium sulphide CuCl Copper (I) chloride CuCl2 Copper (II) chloride Trinity College Dublin, The University of Dublin Polyatomic Ions Ionic Compounds can also contain polyatomic ions Ions which act as discrete units (with an overall charge) Hydroxide Carbonate Bicarbonate Sulfate Sulfite Nitrate Nitrite OH - CO3 2- HCO3 2- SO4 2- SO3 2- NO3- NO2- Most polyatomic ions are oxyanions (oxoanions) Element, usually non-metal, is bonded to 1 or more oxygen atoms Ion with more O atoms takes the nonmetal root and the suffix -ate. Ion with fewer O atoms takes the nonmetal root and the suffix -ite. KNO3 NaNO3 Na2CO3 Potassium Nitrate Sodium Nitrite Sodium Carbonate Trinity College Dublin, The University of Dublin Covalent Compounds Formed between atoms which differ by little (or not at all) in their tendancy to lose or gain electrons (electronegativity) Nucleus of 1 atom attracts valence electrons on another The shared electron pair is localized between atoms. Distinct molecules are formed. Formula is a molecular formula(gives exact number or atoms in each molecule) Trinity College Dublin, The University of Dublin Ionic Compounds Formed between atoms with large difference in their tendancy to lose or gain electrons (electronegativity) Transfer of electron from metal atom to nonmetal Each atom achieves Noble gas configuration Interactions form a 3D ionic solid. Formula is empirical formula (gives ratio of one to another) Trinity College Dublin, The University of Dublin Covalent Bonding Each atom in covalent bond “counts” shared electrons as belonging entirely to itself-to achieve full outer valence An outer-level electron pair not involved in bonding is called a lone pair Bonding pairs represented as lines and lone pairs represented as dots H + xH HxH H H SingleBond xxxxxx Fx F Fx F F Fx SingleBond +x x x x xxx xx xx Asinglebond(mostcommon)containsonebondingpairofelectrons. Asinglebondhasa‘bondorder’of1 Trinity College Dublin, The University of Dublin Double and triple bonds Multiplebondspossible(ofteninvolvingC,N,O) ↑↓ ↑↓ ↑↓ ↑ ↑ Z=8 O:1𝑠2 2𝑠2 2𝑝2 2𝑝1 2𝑝1 𝑥𝑦𝑧 xx O +xOx O xx xx xO xx OO xx Z=7 4 electrons form 2 bonding pairs->double bond. ↑↓ ↑↓ ↑ ↑ ↑ N:1𝑠22𝑠22𝑝2 2𝑝12𝑝1 𝑥𝑦𝑧 6 electrons form 3 bonding pairs->triple bond xx x xx N +xNx NxNx N Nx xxxxx Trinity College Dublin, The University of Dublin Polyatomic structures Methane is a polyatomic molecule with molecular formula CH4 C+xHxHxHxH Total electron pairs available: 8 total valence electrons ÷ 2 = 4 pairs Place the atom with the lowest ionisation energy in the middle and put other atoms symmetrically around it. Use electron pairs to form chemical bonds. If octets aren’t full, make multiple bonds. Carbon has a share in 8electrons, Hydrogens have a share in 2 electrons HH x Hx C xH H C H 4 pairs of electrons 4 C-H single bonds x H H Any remaining pairs of electrons are placed on the central atom. Trinity College Dublin, The University of Dublin Polyatomic structures - ammonia Ammonia is a polyatomic molecule with molecular formula NH3 N+xHxHxH Total electron pairs available: 8 electrons = 4 pairs We place nitrogen in the centre, and use three electron pairs to make three N-H single bonds. The unused electron pair is placed on nitrogen to fill up nitrogen’s octet (because we can’t make multiple bonds). Nitrogen has a share in 8electrons, Hx N xH H N H Hydrogens x 4 pairs of electrons 3 C-H single bonds + 1 lone pair on nitrogen have a share in 2 electrons H H Trinity College Dublin, The University of Dublin 1. Magnesium oxide xx xOx xx xx Mg2+ xOx 2- xx Mg 2. Silicon tetrachloride Si + xCl xCl xCl xCl Cl x Clx SixCl x Cl Cl Example Cl Si Cl Cl Trinity College Dublin, The University of Dublin Moles Mole: amount of substance that contains as many units as there are atoms in 12 g of carbon – 12. This number is called the Avogadro’s 23 number (𝑁 ) = 6.022 x 10 molecules or atoms 1 mol of O = 6.022 x 1023 atoms of oxygen 1 mol of H2O = 6.022 x 1023 molecules of water The mass in grams of one mole of any substance is its molar mass. 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠(𝑚𝑜𝑙) = 𝑚𝑎𝑠𝑠 (𝑔) 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠(𝑛) No. of atoms/molecules/ions = No. of moles x Avogadro’s constant 𝑛𝑜𝑜𝑓𝑎𝑡𝑜𝑚𝑠=𝑛 × 𝑁 𝐴 𝐴 Trinity College Dublin, The University of Dublin Molar mass Eg: Sodium carbonate Na2CO3 2 x Na + 1 x C + 3 x O 2 x 23 + 1 x 12 + 3 x 16 = 106 g/mol Calculate the molar mass of ethanol (C2H5OH) 2 x C + 6 x H + 1 x O = 2 x 12 + 1 x 6 + 1 x 16 = 46 g/mol Trinity College Dublin, The University of Dublin Examples How many moles of Fe are in 5.6 g Fe? How many Fe atoms are contained in the sample? 1 mole of Fe = 56 g 5.6 g Fe = 0.1 mole of Fe. The number of Fe atoms in the sample is 23 22 0.1 mole x 6.022 x 10 atoms/mole = 6.022 x 10 atoms. 5.6 g of iron is not much iron. However, even this small amount contains a huge number of iron atoms. Trinity College Dublin, The University of Dublin Drawing Lewis structures A Lewis Structure is a diagram of an element showing its valence electrons (outer shell electrons) drawn as dots. H Hydrogen Lewis structure We represent pairs of electrons in the same orbital as two dots, unpaired electrons are drawn as isolated, unpaired dots. Oxygen: 8 electrons Valence electrons ↑↓ ↑↓ ↑↓ ↑ ↑ O:1𝑠22𝑠22𝑝2 2𝑝12𝑝1 𝑥𝑦𝑧 An unpaired electron O A pair of electrons 2 pairs of valence electrons 2 unpaired valence electrons Trinity College Dublin, The University of Dublin Lewis structure – Ionic bonds 1. Lithium has 1 outer electron (wants to lose 1, valence = 1), while Bromine 7 outer electrons (wants to gain 1, valence =1): Li Br The formation of Lithium Bromide can be drawn using Lewis Structures (we draw a cross to highlight Lithium’s electron): Lix + Br Li+ xBr- 2. Calcium has 2 valence electrons while chlorine has 7 valence electrons; the formation of calcium chloride is therefore: Cl + xCax + Cl Cl - Ca2+ - xCl x Trinity College Dublin, The University of Dublin Lewis structures – covalent bonds Hydrogen has 1 valence electron but needs 2 to fill its outer shell. By sharing their electrons, two H atoms can fill their outer shells: H+xH HxH A pair of electrons between two atoms represents a bond. We can draw a line instead to represent this bonding electron pair: H+xH HH Fluorine has 7 outer electrons. By sharing their electrons, two F atoms can both have 8 electrons in their outer shell to form F2: FF F Fx F Fx xx xx xx xx +xx xx Trinity College Dublin, The University of Dublin Double and triple bonds Multiplebondspossible(ofteninvolvingC,N,O) ↑↓ ↑↓ ↑↓ ↑ ↑ Z=8 O:1𝑠2 2𝑠2 2𝑝2 2𝑝1 2𝑝1 𝑥𝑦𝑧 xx O +xOx O xx xx xO xx OO xx Z=7 4 electrons form 2 bonding pairs->double bond. ↑↓ ↑↓ ↑ ↑ ↑ N:1𝑠22𝑠22𝑝2 2𝑝12𝑝1 𝑥𝑦𝑧 6 electrons form 3 bonding pairs->triple bond xx x xx N +xNx NxNx N Nx xxxxx Trinity College Dublin, The University of Dublin 1. Magnesium oxide xx xOx xx xx Mg2+ xOx 2- xx Mg 2. Silicon tetrachloride Si + xCl xCl xCl xCl Cl x Clx SixCl x Cl Cl Example Cl Si Cl Cl Trinity College Dublin, The University of Dublin Polyatomic structures Methane is a polyatomic molecule with molecular formula CH4 C+xHxHxHxH Total electron pairs available: 8 total valence electrons ÷ 2 = 4 pairs Place the atom with the lowest ionisation energy in the middle and put other atoms symmetrically around it. Use electron pairs to form chemical bonds. If octets aren’t full, make multiple bonds. Carbon has a share in 8electrons, Hydrogens have a share in 2 electrons HH x Hx C xH H C H 4 pairs of electrons 4 C-H single bonds x H H Any remaining pairs of electrons are placed on the central atom. Trinity College Dublin, The University of Dublin Valence shell electron pair repulsion theory (VSEPR) Trinity College Dublin, The University of Dublin What shapes are these really? Molecules are 3 – dimensional Lewis structures are flat 2D drawings but molecules are 3D. VSEPR theory uses the electron pairs from Lewis Structures to predict the 3D shapes of molecules. A small molecule A large biomolecule DNA Trinity College Dublin, The University of Dublin Why is shape important Properties are highly dependent upon shape. We’ll look at polarities caused by different geometries. There are repercussions of this in almost every property. E.g.: » Boiling / Melting Points (intermolecular forces) » Interactions with other materials and substances » Larger structure/material packing » Biological receptors / enzymes very geometry specific Trinity College Dublin, The University of Dublin VSEPR theory VSEPR = Valence Shell Electron-Pair Repulsion ‘To minimize repulsions, each group of valence electrons around a central atom is located in a way which maximises the distance between the electron pairs.’ H HCH H Trinity College Dublin, The University of Dublin Simple examples Linear molecules (one central atom with two other atoms) such as CO2 have bond angle of 180° to maximise the distance between their bonding pairs: O C O Linear 180 ̊ Trigonal planar molecules (one central atom with three other atoms) such as BF3 have a bond angle of 120° for the same reason: F B Trigonal Planar 120 ̊ F F Trinity College Dublin, The University of Dublin Complex example Phosphorous Pentachloride (PCl5) is one of many molecules that break the octet rule for the central atom. However, we can still determine the 3D shape of the molecule using VSEPR theory: Cl Cl P Cl Cl Trigonal bipyramidal 120 ̊ and 90 ̊ angles Cl Two tetrahedra stacked on top of each other.... Trinity College Dublin, The University of Dublin VSEPR and lone pairs Lone pairs are treated like bonding pairs; they are also repelled by valence electron pairs. However because they are closer to the central atom, lone pairs are slightly more repulsive. This results in a small compression of the bond angles in structures with lone pairs (CH4 = 109°, NH3= 107.5°) HNH H 107.5 ̊ The order of repulsion between different types of electron pairs is as follows: Lone pair – lone pair > lone pair – bond pair > bond pair – bond pair Trinity College Dublin, The University of Dublin Tetrahedral Four electron pairs Electron-pair geometry = tetrahedral Trigonal pyramidal 107.5 ̊ Ammonia, NH3 3 bond pairs 1 lone pair Ammonia has 3 bond pairs and 1 lone pair, so it has a trigonal pyramidal molecular shape Bent Water, H2O 2 bond pairs 2 lone pairs Water has 2 bond pairs and 2 lone pair, so it has a bent or angular molecular shape Effect of lone pair on bond angles 104.5 ̊ 109.5 ̊ Methane, CH4 4 bond pairs No lone pairs Methane has bond pairs, so it has a tetrahedral molecular shape 4 Trinity College Dublin, The University of Dublin VSEPR electron arrangements Electron arrangement is the 3D arrangement of electron pairs around a central atom. As we know, lone pairs and bonding pairs have a slightly different effect on the molecular geometry. We can catalogue the VSEPR predictions for all possible electron arrangements using the following notation: A Xm En En = Number of lone pairs (n) A = element of central atom Xm = element of surrounding atoms (X) and number of atoms (m) Trinity College Dublin, The University of Dublin VSEPR type Bonding Lone pairs (2e-) Geometry Examples pairs AX2 2 0 Linear BeCl2, CaCl2 AX2E2 2 2 Bent H2O, H2S (non – linear) AX3 3 0 Trigonal planar BCl3, BF3 AX3E 3 1 Trigonal NH3 pyramidal AX4 4 0 Tetrahedral CH4 AX5 5 0 Trigonal PCl5 bipyramidal AX6 6 0 Octahedral SF6 Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin H HCH H Methane Bond in the plane of the paper Dashed wedge – bond going backwards into the page Solid wedge – bond coming forward from the page Drawing in 3D Trinity College Dublin, The University of Dublin H N HH Trigonal Pyramidal Drawing in 3D PCl5 Trigonal bipyramid Trinity College Dublin, The University of Dublin Polar bonds vs. Polar molecules A polar covalent bond exists between two atoms of different electronegativities such as C- Cl, C-F, C-Br, C-O, C-N, O-H, N-H, P-Cl, Si-Cl Trinity College Dublin, The University of Dublin Polar bonds vs. Polar molecules A polar covalent bond exists between two atoms of different electronegativities such as C- Cl, C-F, C-Br, C-O, C-N, O-H, N-H, P-Cl, Si-Cl The electrons in the bond are closer to the more electronegative atom, so this becomes δ- (a little bit more negative) while the other end of the bond becomes δ+ ( a little bit more positive) + - C Cl A non-polar covalent bond exists between atoms of the same or close electronegativity such as C-C, C-H, H-H, F-F, Br-Br, Cl-Cl, I-I, O-O Trinity College Dublin, The University of Dublin Polar bonds and non-polar molecule Carbon dioxide (CO2) is a linear molecule by VSEPR theory: OC O Linear 180 ̊ Oxygen is more electronegative than carbon; thus each bond is polarised. However due to the symmetry of the molecule, CO2 is not a polar molecule. δ- δ+ δ- OCO Both ends of CO2 have a partial negative charge, so it’s not polar Trinity College Dublin, The University of Dublin Predicting the polarity of water Oxygen in water (H2O) has to lone pairs, so the bond angle is further reduced to just 104.5° OH H 104.5 ̊ Oxygen is more electronegative than hydrogen, so the bonding electrons are more attracted to oxygen. A partial negative charge builds up around oxygen, and a partial positive charge around hydrogen. As a result, water is a highly polar molecule. δ- O HH δ+ δ+ Trinity College Dublin, The University of Dublin Things you need to understand from L1-4 Structure of the atom. Reading number of protons, calculating number of electrons Understand order of periodic table and the trends Calculating number of moles/molar mass/number of atoms. Know 4 quantum numbers and write electronic configuration Different types of bonding, why they happen, draw lewis structures. VSEPR and Shapes of Molecules Chemical reactions Reactions convert one set of substances to another. The initial substances are called reactants/reagents and the resulting substances are called products. Reactants Products Trinity College Dublin, The University of Dublin Chemical reactions Involve a rearrangement of the atoms in one or more substances. Distinct from physical or nuclear changes e.g. Combustion of methane (CH4), in oxygen (O2) to form carbon dioxide (CO2) and water (H2O) CH4 + O2 → CO2 + H2O Reactants Products Bonds are broken Bonds are formed. All atoms present in reactants must be accounted for in products. This involves balancing the chemical equation. Trinity College Dublin, The University of Dublin Moles Mole: amount of substance that contains as many units as there are atoms in 12 g of carbon – 12. This number is called the Avogadro’s 23 number (𝑁 ) = 6.022 x 10 molecules or atoms 1 mol of O = 6.022 x 1023 atoms of oxygen 1 mol of H2O = 6.022 x 1023 molecules of water The mass in grams of one mole of any substance is its molar mass. 𝑛𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠(𝑚𝑜𝑙) = 𝑚𝑎𝑠𝑠 (𝑔) 𝑚𝑜𝑙𝑎𝑟 𝑚𝑎𝑠𝑠(𝑛) No. of atoms/molecules/ions = No. of moles x Avogadro’s constant 𝑛𝑜𝑜𝑓𝑎𝑡𝑜𝑚𝑠=𝑛 × 𝑁 𝐴 𝐴 Trinity College Dublin, The University of Dublin Applying stoichiometry 2H2 +O2→2H2O The number of moles of each species are contained within the chemical equation (relative stoichiometry). e.g: 1moletoO2 :2molesofH2O Q. How many moles of H2O form from 0.25 moles O2? 1 mole of O2 2 moles of H2O 0.25 moles 0.5 moles of H2O Trinity College Dublin, The University of Dublin Chemical reactions Two laws of import for every reaction you see: 1. Mass Balance 2. Charge Balance 1. The law of conservation of mass: 2H2 +O2→2H2O Mass is neither created nor destroyed in a chemical reaction, it is only converted from one form to another. Number of reactant atoms = Number of product atoms Trinity College Dublin, The University of Dublin Chemical reactions Two laws of import for every reaction you see: 1. Mass Balance 2. Charge Balance 2. The law of charge balance: 2H2 +O2→2H2O If all reactants are neutral, the products will also be neutral overall. Number of reactant charges = Number of product charges Trinity College Dublin, The University of Dublin Chemical equations Tell us what will be formed when reactants combine. Mg + O2→MgO Does this satisfy the law of conservation of mass? No. We have to balance the equation Step 1: Mg 1 Mg Mg 1 Mg 2Mg + + O2 2O O2 2O → MgO 1 Mg 1O → 2MgO 2 Mg 2O Count both sides Add coefficients as needed Count both sides Add coefficients as needed Step 2: Step 3: + 1 O2 → 2MgO Count both sides Balanced! Trinity College Dublin, The University of Dublin Balancing chemical equations Na + H2O→NaOH+ H2 1 Na 2H 1O 1 Na 1O 1H 2H 1Na, 2H, 1O Hydrogens are not balanced, so we need to balance them: Na + 2H2O→2NaOH+ H2 2 Na 2 H 2O 2H 1Na, 4H, 2O But, now sodium are not balanced, we need to balance them: 2Na + 2H2O → 2NaOH + H2 1 Na 4 H 2O 1Na, 3H, 1O 2Na, 4H, 2O 2Na, 4H, 2O 2Na, 4H, 2O Trinity College Dublin, The University of Dublin Reaction components Why is that important? The states of reactants and products are important. They inform on whether a reaction is: » Feasible » Likely going to be fast or slow ▪ (two solids at RT? Probably not) ▪ (homogeneity vs. heterogeneity) » Actually taking place ▪ (ionic dissociation) » Being driven by a particular effect ▪ (precipitation, gas evolution) Trinity College Dublin, The University of Dublin State symbols: Reaction components » (s) = solid » (l) = liquid » (g) = gaseous » (aq) = aqueous solution Dissolved in water Is a change: physical or chemical or both? State symbols will show this Dissolution Homogeneous Water reaction splitting. Expected in all balanced reaction equations. Trinity College Dublin, The University of Dublin State symbols: » (s) = solid » (l) = liquid » (g) = gaseous » (aq) = aqueous solution Most metal and non-metal elements in their standard states. Ionic compounds (e.g. NaCl, BaCl2, TiO2) Water, alcohol, mercury H2, O2, He, N2, N2O, CO, CO2, Noble gases, methane (CH4), ammonia (NH3) Reaction components Trinity College Dublin, The University of Dublin Reaction components 2Na + 2H2O → 2NaOH + H2 2Na(s) + 2H2O(l)→2NaOH(aq)+ H2(g) HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (aq) Trinity College Dublin, The University of Dublin Various types of solutions Example State of State of solute State of solution solvent Air, natural gas Gas Gas Gas Antifreeze Liquid Liquid Liquid Brass Solid Solid Solid Carbonated water Liquid Gas Liquid Seawater, sugar solution Liquid Solid Liquid Hydrogen in platinum Solid Gas Solid Trinity College Dublin, The University of Dublin Water: the common solvent Water is one of the most important substances on the earth. A H2O molecule is bent or V-shaped with an H-O-H angle of 104.5 o Water is a polar molecule. This polarity gives water its ability to dissolve compounds. OH H 104.5 ̊ δ- O HH δ+ δ+ Trinity College Dublin, The University of Dublin Hydration and Dissolution δ+ parts of water molecules attracted to negatively charged anions δ- part attracted to the positively charged cations. Hydration of ions tends causes salt to dissolve. Strong forces (between positive and negative ions in solid) replaced by strong water-ion interactions. Trinity College Dublin, The University of Dublin Solubility Ability of a given substance (solute), to dissolve (in a solvent) Measured in terms of the maximum amount of solute dissolved at equilibrium. Trinity College Dublin, The University of Dublin 1. Ionic substances Solubility Solubility of ionic substances in water varies. NaCl is quite soluble in water, whereas AgCl very slightly soluble Solubility in water depends on – the relative attractions of the ions for each other – attractions of the ions for water molecules(hydration) Trinity College Dublin, The University of Dublin Solubility Water also dissolves non-ionic substances e.g. ethanol (C2H5OH) Ethanol contains a polar O-H bond like those in water. In general, polar substances are only soluble in polar solutions. 2. Non – ionic substances Trinity College Dublin, The University of Dublin Levels of solubility Very soluble Freely soluble Soluble Sparingly soluble Slightly soluble Very slightly soluble Insoluble Trinity College Dublin, The University of Dublin Concentration Ratio of solute to solvent Can be expressed in various units depending on the circumstances Sometimes units of grams per dm3 (or grams per litre) e.g. a solution containing 20 g of sodium chloride dissolved in 1 dm3 3 -3 of solution has a concentration of 20 g/dm (20 g dm ) We are more interested in the chemical amount of substance present rather than the mass So we again return to moles This time as mol/L or Molarity (M). Trinity College Dublin, The University of Dublin Molarity Number of moles dissolved per Litre of solution. The units are = mol/dm3 or mol/L. 𝑀𝑜𝑙𝑎𝑟𝑖𝑡𝑦= 𝑚𝑜𝑙𝑒𝑠𝑜𝑓𝑠𝑜𝑙𝑢𝑡𝑒 𝑙𝑖𝑡𝑟𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 𝑁𝑢𝑚𝑏𝑒𝑟 𝑜𝑓 𝑚𝑜𝑙𝑒𝑠(𝑛) = 𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛 × 𝑣𝑜𝑙𝑢𝑚𝑒 mol Mol L-1 L (1dm3 =1L) n cV Trinity College Dublin, The University of Dublin 1. Percent composition 𝐴𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝐴𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 Can be divided into: a) Percentage composition by mass (w/w %) b) Percentage composition by volume (v/v %) c) Percentage composition by mass/volume (w/v %) × 100 Trinity College Dublin, The University of Dublin a) Percent composition: w/w % Let's consider a 12% by weight sodium chloride solution. Such a solution would have 12 g of sodium chloride for every 100 g of solution. To make such a solution, you could weigh out 12 g of sodium chloride, and then add 88 g of water, so that the total mass for the solution is 100 g. 𝐺𝑟𝑎𝑚𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝐺𝑟𝑎𝑚𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ×100 Trinity College Dublin, The University of Dublin b) Percent composition: v/v % Used whenever a solution is prepared by mixing pure liquid solutions 70% v/v rubbing alcohol may be prepared by taking 70 ml of isopropyl alcohol and adding sufficient water to obtain 100 ml of solution. Solutions made to a specific volume percent concentration typically are prepared using a volumetric flask. 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 ×100 Trinity College Dublin, The University of Dublin c) Percent composition: w/v % amount of solute in grams but amount of solution in millilitres. Physiologic or isotonic saline is a 0.9% aqueous solution of NaCl. 0.9% saline = 0.9 g of NaCl diluted to 100 mL of deionized water, where NaCl is the solute and deionized water is the solvent. 𝑊𝑒𝑖𝑔h𝑡 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 (𝑔) × 100 𝑉𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 (𝑚𝐿) Trinity College Dublin, The University of Dublin Example: A 4 g sugar cube (Sucrose: C12H22O11) is dissolved in a 350 ml teacup of 80 °C water. What is the percent composition by mass of the sugar solution? Given: Density of water at 80 °C = 0.975 g/ml Step 1 - Determine mass of solute We were given the mass of the solute in the problem. The solute is the sugar cube. masssolute = 4 g Step 2 - Determine mass of solvent mass = density x volume = 0.975 g/ml x 350 ml masssolvent = 341.25 g Trinity College Dublin, The University of Dublin Example: Step 3 - Determine the total mass of the solution msolution = msolute + msolvent msolution = 4 g + 341.25 g msolution = 345.25 g Step 4 - Determine percent composition by mass of the sugar solution. Percent composition = 𝑚𝑠𝑜𝑙𝑢𝑡𝑒 × 100% 𝑚𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 = 4 ×100% 345.25 = 1.16% Dilution To save time and space in the lab, routinely used solutions are often purchased or prepared in concentrated form called stock solutions. Water is then added to achieve the desired molarity for a particular solution. This process is called dilution. Trinity College Dublin, The University of Dublin Dilution When calculating the dilution factors, it is important to remember the units of volume and concentration remain constant. 𝑴𝟏𝑽𝟏 = 𝑴𝟐𝑽𝟐 𝑀𝑜𝑟𝑀 =𝑐𝑜𝑛𝑐𝑒𝑛𝑡𝑟𝑎𝑡𝑖𝑜𝑛𝑜𝑓𝑡h𝑒𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 12 𝑉𝑜𝑟𝑉 =𝑣𝑜𝑙𝑢𝑚𝑒𝑜𝑓𝑡h𝑒𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 12 For example, what volume of 16 M sulphuric acid must be used to prepare 1.5 L of a 0.1 M sulphuric acid solution? 𝑀 =16𝑀 𝑉 =? 11 𝑀 =0.1𝑀 𝑉 =1.5𝐿 22 𝑀𝑉=𝑀𝑉 11 22 16Mx𝑉 =0.1Mx1.5L 1 -3 𝑉 =9.4x10 Lor9.4mL 1 Trinity College Dublin, The University of Dublin Serial dilutions Serial dilutions involve diluting a stock or standard solution multiple times in a row. Typically, the dilution factor remains constant for each dilution, resulting in an exponential decrease in concentration. For example, a ten-fold serial dilution could result in the following concentrations: 1 M, 0.1 M, 0.01 M, 0.001 M, and so on. Trinity College Dublin, The University of Dublin Dilution of original culture Trinity College Dublin, The University of Dublin The concept of Limiting Reagents Let’s start with an example: Limiting Reagent Trinity College Dublin, The University of Dublin The concept of Limiting Reagents When molecules react to form products, we also need to consider this N2 (g) + 3 H2 (g)→2 NH3 (g) H2 N2 For the reaction to go to completion, we need to remember that each N2 requires 3 H2 molecules to form 2 NH3. Trinity College Dublin, The University of Dublin Stoichiometric Mixtures H2 N2 NH3 3:1 Relative amounts that match the numbers in the balanced equation: N2 (g) + 3 H2 (g)→2 NH3 (g) This type of mixture is called a stoichiometric mixture All reactants will be consumed to form products. Before the reaction After the reaction Trinity College Dublin, The University of Dublin Limiting Reagents N2 (g) + 3 H2 (g)→2 NH3 (g) Remember that each N2 still requires 3 H2 molecules to form 2 moles of NH3 H2 N2 NH3 The H2 molecules are used up before all the N2 molecules are consumed. Amount of hydrogen limits the amount of product NH3 that can form. H2 is the limiting reagent. Before the reaction After the reaction Trinity College Dublin, The University of Dublin Determination of limiting reagent 1. Using reactant quantities e.g 25 kg of N2 and 5 kg of H2 are mixed to form ammonia. Calculate the mass of ammonia produced when the reaction is run to completion (until one of the reactants is completely consumed)? 1. Get balanced equation 2N2H 1N3H N2 (g) + H2 (g)→NH3 (g) Count both sides Add coefficients as needed 2 N, 2 H N (g) + H (g)→2NH (g) 223 Count both sides Add coefficients as needed Count both sides Balanced! 1 N, 3 H N2 (g) + 3H2 (g)→2NH3 (g) Trinity College Dublin, The University of Dublin 2. Calculate number of moles of each reactant ForN2:No.ofmoles= 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 ForH2:No.ofmoles= 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 =25×1000𝑔 =893molN2 28 𝑔Τ𝑚𝑜𝑙 =5×1000𝑔 =2480molH2 2 𝑔Τ𝑚𝑜𝑙 3. Compare the mole ratio of the substances required in balanced equation with the mole ratio of reactants actually present Inthebalancedequationwehave:3𝑚𝑜𝑙𝐻2 =3 1𝑚𝑜𝑙𝑁2 Intheexperimentwehave:2480𝑚𝑜𝑙𝐻2 =2.78 893 𝑚𝑜𝑙 𝑁2 mole ratio of H2 to N2 is smaller than required-> H2 must be the limiting reagent Trinity College Dublin, The University of Dublin 4. To calculate the amount of product (ammonia) formed, we must use the amount of hydrogen. N2 (g) + 3H2 (g)→2NH3 (g) From balanced equation: 3 moles of H2 give 2 moles of NH3 3 moles H2 = 2 moles NH3 (we actually have 2480 mol H2) 2480 moles H2 = X =2480 ×2𝑚𝑜𝑙𝑁𝐻3 =1650molNH3 3𝑚𝑜𝑙𝐻2 Convert the moles to mass: Mass = moles x molecular weight = 1650 mol x 17 g/mol = 28 kg NH3 Trinity College Dublin, The University of Dublin N2 + 3 H2 → 2 NH3 Moles from balanced eqn. 132 Mass 25 kg 5 kg Calculated moles = 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 = 25 ×1000𝑔 28 𝑔Τ𝑚𝑜𝑙 = 893 mol N2 5 ×1000𝑔 2 𝑔Τ𝑚𝑜𝑙 = 2480 mol H2 Actual ratio 3 𝑚𝑜𝑙 𝐻2 = 3 1 𝑚𝑜𝑙 𝑁2 Calculated ratio 2480 𝑚𝑜𝑙 𝐻2 = 2.78 893 𝑚𝑜𝑙 𝑁2 Moles of ammonia 2480 ×2𝑚𝑜𝑙𝑁𝐻3 3 𝑚𝑜𝑙 𝐻2 = 1650 mol NH3 Mass of ammonia 1650 mol x 17 g/mol = 28 kg NH3 Trinity College Dublin, The University of Dublin Determination of limiting reagent 2. Using quantities of products formed Determination of limiting reagent by calculating the amounts of products formed by completely consuming each reactant. The reactant that produces the smallest amount of product must run out first and thus will be the limiting reagent. Consider previous example: 25kgofN2 and5kgofH2 1. Get balanced equation N2 (g) + 3H2 (g)→2NH3 (g) Trinity College Dublin, The University of Dublin 2. Calculate number of moles of each reactant ForN2:No.ofmoles= 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 ForH2:No.ofmoles= 𝑚𝑎𝑠𝑠 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 =25×1000𝑔 =893molN2 28 𝑔Τ𝑚𝑜𝑙 =5×1000𝑔 =2480molH2 2 𝑔Τ𝑚𝑜𝑙 N2 (g) + 3H2 (g)→2NH3 (g) 3. Calculate moles of products formed by completely consuming each reactant. 1 mole of N2 gives 2 moles of NH3 1 mole H2 = 2 moles NH3 (we actually have 893 mol H2) 893 moles N2 = X 893molN2 ×2𝑚𝑜𝑙𝑁𝐻3 =1790molNH3 1 𝑚𝑜𝑙 𝑁2 Trinity College Dublin, The University of Dublin 3 moles of H2 give 2 moles of NH3 3 moles H2 = 2 moles NH3 (we actually have 2480 mol H2) 2480 moles H2 = X =2480mol𝐻2×2𝑚𝑜𝑙𝑁𝐻3 =1650molNH3 3𝑚𝑜𝑙𝐻2 If all the N2 reacted we would form 1790 mol of NH3 If all the H2 reacted we would form 1650 mol NH3 H2 must be limiting agent Convert the moles to mass: Mass = moles x molecular weight = 1650 mol x 17 g/mol = 28 kg NH3 Trinity College Dublin, The University of Dublin Example: Nitrogen gas can be prepared by passing ammonia over solid copper (II) oxide at high temperatures. The other products of the reaction are solid copper and water vapour. If a sample containing 18.1 g of NH3 is reacted with 90.4 g of CuO, which is the limiting reagent? How many grams of N2 will be formed? Reaction (unbalanced): NH3 (g)+CuO(s)→N2 (g)+Cu(s)+H2O(g) Trinity College Dublin, The University of Dublin 1. Get balanced equation NH3 (g)+CuO(s)→N2 (g)+Cu(s)+H2O(g) 1N 1Cu 2N 1Cu 2H 3H1O1O 2NH3 (g) + CuO (s)→N2 (g) + Cu (s) + H2O (g) 2NH3 (g) + CuO (s)→N2 (g) + Cu (s) + 3H2O (g) 2NH3 (g) + 3CuO (s)→N2 (g) + 3Cu (s) + 3H2O (g) Count both sides Add coefficients as needed Count both sides Balanced! Trinity College Dublin, The University of Dublin 2. Calculate number of moles of each reactant For NH3: No. of moles = 𝑚𝑎𝑠𝑠 = 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 For CuO: No. of moles = 𝑚𝑎𝑠𝑠 = 𝑚𝑜𝑙𝑒𝑐𝑢𝑙𝑎𝑟 𝑤𝑒𝑖𝑔h𝑡 18.1 𝑔 = 1.06 mol NH3 17 𝑔Τ𝑚𝑜𝑙 3. Calculate moles of products formed by completely consuming each reactant. (Second method) 1.06molNH3× 1𝑚𝑜𝑙𝑁2 =0.530molN2 2 𝑚𝑜𝑙 𝑁𝐻3 1.14molCuO× 1𝑚𝑜𝑙𝑁2 =0.380molN2 3 𝑚𝑜𝑙 𝐶𝑢𝑂 Because a smaller amount of N2 is produced from CuO than from NH3, the limiting reagent is CuO. 90.4 𝑔 79.55 𝑔Τ𝑚𝑜𝑙 = 1.14 mol CuO Trinity College Dublin, The University of Dublin 4.Calculate the mass of product formed. Because CuO is the limiting reagent, we must used the amount of CuO to calculate the amount of N2 formed. Convert the moles to mass: Mass of N2 = moles x molecular weight = 0.380 mol x 28 g/mol = 10.6 g N2 Trinity College Dublin, The University of Dublin Theoretical and Percentage Yields The amount of a product formed when the limiting reagent is completely consumed is called the theoretical yield. In the above example, the amount of nitrogen formed (10.6 g) represents the theoretical yield. This is the maximum amount of nitrogen that can be produced from the quantities of reactants used. The theoretical yield is seldom obtained (side reactions that involve one ore more of the reactants or products) The actual yield of product is often given as a percentage of the theoretical yield. This is called as the percentage yield. Trinity College Dublin, The University of Dublin Theoretical and Percentage Yields 𝑃𝑒𝑟𝑐𝑒𝑛𝑡𝑎𝑔𝑒 𝑦𝑖𝑒𝑙𝑑 = 𝐴𝑐𝑡𝑢𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 × 100% 𝑇h𝑒𝑜𝑟𝑒𝑡𝑖𝑐𝑎𝑙 𝑦𝑖𝑒𝑙𝑑 In the above example, if the reaction considered actually gave 6.63 g of nitrogen instead of predicted 10.6 g, the percentage yield will be: 6.63 × 100 = 62.5% 10.6 Trinity College Dublin, The University of Dublin Problem solving strategy 1. Write and balance the equation. 2. Convert the known masses of substances to moles. (don’t forget to do any unit conversions if there is any). 3. Determine the limiting reagent. 4. Using the amount of limiting reagent and the appropriate mole ratios, compute the number of moles of the desired product. 5. Convert from moles to grams, using the molar mass. Trinity College Dublin, The University of Dublin Outline Limiting reagents Theoretical and percentage yields Types of chemical reactions Trinity College Dublin, The University of Dublin Reaction components An electrolyte is a substance that dissolves to give a solution that contains ions. The ions allow the solution to conduct electricity. 1. A strong electrolyte is a compound that dissolves to give a solution that contains mainly ions. NaCl→Na+(aq)+Cl- (aq) An ionic solution that conducts electricity Trinity College Dublin, The University of Dublin Electrolytes A weak electrolyte is a compound that dissolves to give a solution that contains mostly molecules. won’t conduct CH3COOH (aq) → H+ + CH3COO- CH3COOH (aq) ⇌ H+ + CH3COO- A non-electrolyte is a substance that dissolves to solution but doesn’t contain ions and therefore, doesn’t conduct electricity. E.g. Glucose (C6H12O6), Ethanol (C2H5OH) Trinity College Dublin, The University of Dublin Types of chemical reactions Types of solution reactions: Precipitation reactions Acid – base reactions Oxidation – reduction reactions Trinity College Dublin, The University of Dublin 1. Precipitation reactions A precipitation reaction usually occurs when solutions of 2 strong electrolytes are mixed. NaCl (aq) + AgNO3 (aq)→ AgCl (s) + NaNO3 (aq) AgNO3 (aq) NaCl (aq) AgCl (s) + NaNO3 (aq) Trinity College Dublin, The University of Dublin 1. Precipitation reactions Precipitation is the process in which a solute comes out of solution rapidly as a solid, called a precipitate. Precipitate Trinity College Dublin, The University of Dublin 2. Acid – base reactions An acid is a species which donates a proton. A base is a species which accepts a proton. Neutralization is the reaction of an H+ (H3O+) ion from the acid and the OH- ion from the base to form water, H2O. HCl (aq) + NaOH (aq)→NaCl (aq) + H2O (aq) acid base salt water Redox in biology Living cells utilise redox chemistry to generate energy. Respiration is the oxidation of sugar to form CO2 and H2O: C6H12O6(s)+6O2 (g)→6CO2(g)+6H2O(g) Oxygen is a strong oxidising agent, meaning it can accept electrons easily. Oxygen acts as a final acceptor of electrons in the electron transport chain, one method mitochondria (the 'power plant' of cells) use to produce ATP (a biological form of energy). Biological systems usually transfer electrons as a chemical bond between a pair of atoms; very often between a pair of H atoms. Trinity College Dublin, The University of Dublin Vitamin B: NAD and FAD NAD (Nicotinamide Adenine Dinucleotide) is derived from Niacin (Vitamin B3) while FAD ( Flavin Adenine Dinucleotide) is derived from Riboflavin (Vitamin B2), both of these co- enzymes carry energy during respiration in the form of hydrogen: FAD (oxidised) FADH2 (reduced) NAD (oxidised) NADH + H+ (reduced) Trinity College Dublin, The University of Dublin NAD can be reversibly reduced. They act as coenzymes in many biological processes. NAD+ gets reduced to NADH in glycolysis, among other processes. NADH gets oxidised in anaerobic metabolism. NAD Vitamin B: NAD and FAD NAD+ NADH (Oxidized) (Reduced) Trinity College Dublin, The University of Dublin Redox reactions Redox Oxidation number Rules for Oxidation number Trinity College Dublin, The University of Dublin Na: 11 electrons Cl: 17 electrons Chloride ion (Cl-) Chlorine gains 1 electron Reduction reaction Redox reactions Sodium ion (Na+) Sodium loses 1 electron Oxidation reaction Trinity College Dublin, The University of Dublin Redox reactions Redox reactions involve reduction or oxidation: Oxidation is the loss of electrons. Reduction is the gain of electrons. OILRIG: Oxidation is loss, Reduction is gain. e.g. 3Ag+ (aq) + Al (s) → 3Ag (s) + Al3+ (aq) + Eachsilvercationhasgained1electron→Ag isreduced Each aluminium atom has lost 3 electrons →Al is oxidised Oxidation Numbers: An arbitrary system to allow us to keep track of electrons; Electrons are redistributed as elements react with each other. Trinity College Dublin, The University of Dublin Example redox Na (s) + Cl2 (g) → Na+Cl- (s) Charge Na+ Cl2 Na1+ Cl20 Stoichiometry Trinity College Dublin, The University of Dublin Oxidation numbers: Rules 1. Oxidation number (O.N.) of an ion equals the charge on that ion. Mg2+ O.N. = +2 Cl- O.N. = -1 O2- O.N. = -2 Al3+ O.N. = +3 2. The O.N. of an uncharged (non-ionised) element is zero. Mg (s) O.N. = 0 O2 (g) O.N. = 0 Trinity College Dublin, The University of Dublin Oxidation numbers: Rules 3. An increase in O.N. indicates oxidation. O.N. = 0 O.N. = +2 Mg–2e- → Mg2+ (O.N. 4. A decrease in O.N. indicates reduction. O.N. = 0 O.N. = -1 0→2) Cl2 + 2e- → 2Cl- (O.N. 0 →-1) Special case: In some reactions that produce multiple products you may see a species is both reduced and oxidised: Disproportionation reaction H2O2 (aq)→2 H2O (l) + O2 (g) O.N. = -1 O.N. = -2 O.N. = 0 Trinity College Dublin, The University of Dublin Oxidation numbers: Rules 5. O.N. of Hydrogen: » -1 in compounds with metals » +1 in compounds with non-metals O.N. of Oxygen: » Always -2 » except in peroxide, O22- where each oxygen atom has an O.N. of -1. O.N. of Fluorine: -1 in all its compounds O.N. of other Halogens: -1, except if bound to oxygen or a halogen with larger atomic weight. Trinity College Dublin, The University of Dublin Oxidation numbers: Rules 6. The O.N. of a compound with several atoms is equal to the sum of the O.N. of all the atoms within that species. Eg 1. H2= H + H = 0 + 0 O.N. = 0 Eg2. NaCl=Na++Cl- =(+1)+(-1)=0 O.N.=0 Trinity College Dublin, The University of Dublin Examples What is the Oxidation Number of Sulfur in: A) SO2 and B) SO42-? A) SO2 Sum of oxidation numbers of S and 2(O) equals zero (rule 6) [O.N. of Sulfur] + [2 x (O.N. of Oxygen)] = 0 x + 2(-2) =0 (rule5...O.N.ofOis-2) x =+4 So, the O.N. of Sulfur in SO2 is +4 Trinity College Dublin, The University of Dublin Examples What is the Oxidation Number of Sulfur in: A) SO2 and B) SO42-? B) SO42- Sum of all oxidation numbers = -2 (rule 1 and 6) [O.N. of Sulfur] + [4 x (O.N. of Oxygen)] = -2 x + 4(-2) =-2 (rule5...O.N.ofOis-2) x =+6 So, the O.N. of Sulfur in SO2 is +6 Trinity College Dublin, The University of Dublin Examples What is the O.N. of sulfur in hydrogen sulfide (H2S)? 2H+S = 0(rule1) Hydrogen combined with non-metal has O.N. = +1 (rule 5) 2(+1) + x = 0 x = -2 So, the O.N. of Sulfur in H2S is -2 Trinity College Dublin, The University of Dublin More examples... What is the O.N. of nitrogen in the nitrate ion (NO3-)? »N+3(O) =-1 »x +[3x(-2)]=-1 »x +(-6) =-1 » x=-1+6 =+5 O.N.ofN=+5 What is the O.N. of chlorine in HClO3? »H +Cl+3(O) =0 »+1+ x +[3x(-2)]=0 »1+x-6 =0 » x=6-1=5 O.N.ofCl=+5 Trinity College Dublin, The University of Dublin Oxidising and reducing agents Oxidising Agent is a species which causes oxidation; it is reduced during a redox reaction. Nomenclature: [O] Reducing Agent is a species which causes reduction; it is oxidised during a redox reaction. Nomenclature: [H] Cr2O72- (aq) + Fe2+ (aq) + H+ (aq) → Fe3+ (aq) + Cr3+ (aq) + H2O(l) Trinity College Dublin, The University of Dublin balance the equation before finding the oxidation numbers Cr2O72- (aq) + 6Fe2+ (aq) + 14H+ (aq) → 6Fe3+ (aq) + 2Cr3+ (aq) + 7H2O(l) Cr2O72- (aq) + 6Fe2+ (aq) + 14H+ (aq) → 6Fe3+ (aq) + 2Cr3+ (aq) + 7H2O(l) +6 -2 +2 +1 +3 +3 +1 -2 Working out the oxidation numbers (in red) for the above reaction we see that Cr is reduced (gains e-) while Fe is oxidised (loses e-). Therefore: Cr2O72- (Dichromate) is the oxidising agent and is reduced in the reaction. Fe2+ (Iron (II)) is the reducing agent and is oxidised in the reaction. Trinity College Dublin, The University of Dublin Balancing in redox – charge balance In a chemical reaction, electrons cannot be created nor destroyed, they are transferred from one species to another. Because electrons are charged, the total charge of reactants must equal the total charge of the products. You are given the unbalanced redox equation: Al(s)+H+ (aq)→Al3+(aq)+H2 (g) 1 Al 1 H 1 Al 2 H Balance the number of atoms on either side of the equation: Al(s)+2H+ (aq)→Al3+(aq)+H2 (g) 1 Al 2 H 1 Al 2 H Trinity College Dublin, The University of Dublin Balancing in redox – charge balance Balance charges on both sides of the equation: Al(s)+2H+ (aq)→Al3+(aq)+H2 (g) 0 +2 +3 0 The left side has charge +2, and the right side has charge +3. Lowest Common Multiple = 6 So we need a +6 charge on both sides to balance the equation. We can obtain this by multiplying all H's by 3 and all Al's by 2. Fully balanced equation is: 2Al (s) + 6H+ (aq) → 2Al3+ (aq) + 3H2 (g) 0 +6 +6 0 Trinity College Dublin, The University of Dublin Balancing in redox II – charge balance Sn (s) + Fe3+ (aq) → Sn2+ (aq) + Fe2+ (s) 0 +3 +2 +3 +4 Left species has charge +3 and Right species have charge +2 each Lowest Common Multiple =6 So we need a 6+ charge on both sides of the equation. We can obtain this by multiplying each Fe species by 2. Sn (s) + 2Fe3+ (aq) → Sn2+ (aq) + 2Fe2+ (s) The equation is now balanced with a charge of +6 on both sides. +2 Trinity College Dublin, The University of Dublin Balancing in redox II – charge balance Sn (s) + 2Fe3+ (aq) 0 2 x 3 = +6 +6 → Sn2+ (aq) +2 + 2Fe2+ (s) 2 x 2 = +4 The equation is now balanced with a charge of +6 on both sides. +6 Trinity College Dublin, The University of Dublin Summary Oxidizing agent Reducing agents No of electrons Gained Lost Oxidation state Decreases Increases Substance Reduced Oxidized Oxidizers gain electrons They undergo reduction => they are reduced Reducing agents lose electrons They undergo oxidation => are oxidized Oxidation is the loss of electrons. Reduction is the gain of electrons. Trinity College Dublin, The University of Dublin Half-reactions Redox reactions can be split into reduction and oxidation half-reactions. Chemists use half- reactions to make it easier to see the electron transfer, and it also helps when balancing redox reactions. Cu2+ + Mg→Cu + Mg2+ Split the ionic equation into 2 parts in terms of Cu and Mg: Mg→Mg2+ + 2e- Cu2+ +2e- → Cu Electron half- equations Trinity College Dublin, The University of Dublin Half-reactions Working out electron-half-equations and using them to build ionic equations Chlorine gas oxidizes iron(II) ions to iron(III) ions. In the process, the chlorine is reduced to chloride ions. From this information, the overall reaction can be obtained. Step 1: chlorine gas is reduced to chloride ions Cl2 →Cl- But it is not balanced: Cl2 →2 Cl- Trinity College Dublin, The University of Dublin Half-reactions Working out electron-half-equations and using them to build ionic equations Step 2: To completely balance a half-equation, all charges and extra atoms must be equal on the reactant and product sides. In order to accomplish this, the following can be added to the equation: electrons water hydrogen ions (unless the reaction is being done under alkaline conditions, in which case, hydroxide ions must be added and balanced with water) Trinity College Dublin, The University of Dublin Half-reactions Working out electron-half-equations and using them to build ionic equations Thus the fully balanced half-reaction is: Cl2 +2e-→2Cl- Next, lets consider the iron half-reaction. Iron (II) is oxidised to iron (III): Fe2+ → Fe3+ The atoms are balanced, but the charges are not. Fe2+→Fe3+ + e- Trinity College Dublin, The University of Dublin Half-reactions Working out electron-half-equations and using them to build ionic equations Step 3: combination of two half-equations This reaction has 2 electrons Cl2 + 2e- → 2Cl- Fe2+→Fe3+ + e- 1 x ( Cl2 + 2e-→2Cl-) 2 x (Fe2+→Fe3+ + e-) But this has only 1 electron Cl2 + 2e- + 2 Fe2+→2Cl- + 2Fe3+ + 2e- Cl2 + 2 Fe2+→2Cl- + 2Fe3+ Recap about atoms Atoms are composed of protons, neutrons and electrons) "!X Examples: !𝐻 #"He Openstax.org A = mass number Z = atomic number (no of neutrons = A – Z) X = element symbol Particle Symbol Charge Mass (x 10-27 (x 10-19 Coul.) kg) Proton p +1.60218 1.672623 Neutron n 0 1.674929 Electron e -1.60218 0.000911 Trinity College Dublin, The University of Dublin Isotopes Atoms of an element that have different numbers of neutrons and therefore different mass numbers. Most elements occur in nature in a particular isotopic composition. Example: Ø All neutral carbon atoms (Z = 6) have 6 protons and 6 electrons, Ø 98.89% of naturally occurring carbon atoms have 6 neutrons (A = 12) Ø A small percentage (1.11%) have 7 neutrons (A = 13), Ø Less than 0.01% have 8 (A = 14). 3 naturally occurring isotopes of carbon: "#C , "$C , and "%C !!! Trinity College Dublin, The University of Dublin Example: 1 An isotope of molybdenum has 54 neutrons. What is its atomic symbol? Molybdenum has an atomic number of 42-> no of protons = 42 Mass number = number of protons + no of neutrons = 42 + 54 = 96 Thus the atomic symbol is $%Mo #" Trinity College Dublin, The University of Dublin Example: 2 An element consists of 1.40% of an isotope with mass 203.973 u, 24.10% of an isotope with mass 205.9745 u, 22.10% of an isotope with mass 206.9759 u and 52.40% of an isotope with mass 207.9766 u. Calculate the average atomic mass and identify the element. 1.4×203.973 + 24.1×205.9745 + 22.10×206.9759 +(52.4×207.9766) 100 = 207.2 Therefore, the element is Pb. Trinity College Dublin, The University of Dublin Chemical versus Nuclear Reactions Chemical Reactions 1. One substance converted to another, 1. atoms don’t change identity. 2. Orbital electrons involved, bonds 2. break and form; nucleus is not involved Nuclear Reactions Atoms of one element typically change into atoms of another Protons, neutrons and other particles are involved; orbital electrons rarely take part. Reactions accompanied by large changes in energy and measurable changes in mass. Rates affected by number of nuclei, not by temperature or catalysts. 3. Relatively small changes in energy 4. Reaction rates influenced temperature, concentration, catalysts. 3. 4. by Trinity College Dublin, The University of Dublin Radioactive decay Unstable nucleus emits radiation, forming another nucleus (and producing one or more particles) )*𝐶à)*𝑁+ -𝑒 ( + ,) A and Z must be conserved. Z values must give the same sum on the both sides of the equation (6 = 7 -1 ), as must the A values (14 = 14 + 0). Trinity College Dublin, The University of Dublin 1. Alpha decay emission – heavy elements Helium nucleus ejected, leaving daughter nucleus Ø Note difference in A and Z Ø Note A and Z balanced on both sides Eg. α-decay of uranium-238 #$5𝑈 → #$%𝑇h + %𝐻𝑒 4# 46 # 238Uà234Th + %#𝛼 𝑡"/# = 4.48×104 𝑦𝑒𝑎𝑟𝑠 210Po à 206Pb + %#𝛼 𝑡"/# = 138 𝑑𝑎𝑦𝑠 Unstable nuclides beyond bismuth (Z = 83). Trinity College Dublin, The University of Dublin (which is ejected) 90Sr􏰍90Y + 􏰍􏰍 + Energy 14C􏰍14N + 􏰍􏰍 + Energy 􏰍􏰍􏰍􏰍 􏰍 􏰍􏰍 􏰍􏰍􏰍􏰍􏰍 􏰍􏰍􏰍􏰍 􏰍 􏰍􏰍􏰍􏰍 􏰍􏰍􏰍􏰍􏰍 􏰍􏰍􏰍􏰍 􏰍 􏰍 􏰍􏰍􏰍􏰍 neutrino). Common for proton deficient nuclides, high n/p ratio Ø ( β–particle or 86"𝑒 ) 131I 􏰍 131Xe + 􏰍􏰍 + Energy 2. Beta decay 􏰍􏰍􏰍 􏰍 􏰍􏰍􏰍 􏰍 􏰍􏰍􏰍􏰍 􏰍 􏰍􏰍 Neutron converted to proton (which remains), and β–particle ØNote Z chTarininty Cgolleege Dsub,linA, The Udnivoersiety osf Dnubli’nt (remember mass of electron) "%𝐶 → "%𝑁 + 6𝑒 ! 9 8" 90Sr à 90Y + β8 + Energy 14C à 14N + β8 + Energy 𝑡"/# = 30 𝑦𝑒𝑎𝑟𝑠 𝑡"/# = 5730 𝑦𝑒𝑎𝑟𝑠 𝑡"/# = 8 𝑑𝑎𝑦𝑠 131I à 131Xe + β8 + Energy Common for proton deficient nuclides, high n/p ratio Trinity College Dublin, The University of Dublin electromagnetic radiation, as gamma (􏰍)-ray 3. 𝛾 – emission 􏰍􏰍􏰍􏰍􏰍 􏰍 􏰍􏰍􏰍􏰍 􏰍 􏰍 Nucleons reorganise into more stable arrangement. Normally happens along with other radiation events High energy photons released. Trinity College Dublin, The University of Dublin Ø-𝛾 or γ Note, no change in atomic number or mass 44:𝑇𝑐 → 44𝑇𝑐 + 𝛾 %$ %$ 􏰍􏰍 􏰍􏰍 Trinity College Dublin, The University of Dublin Most common forms of emission-Recap α β γ Charge +2 -1 0 Mass 6.64 x 10-24 9.11 x 10-28 - Penetrating Power (relative) 1 100 10000 Nature of Radiation "!𝐻𝑒 Nuclei Electrons High Energy Photons Trinity College Dublin, The University of Dublin Trinity College Dublin, The University of Dublin 4. Positron Emission Unstable nuclides which are neutron deficient. Proton converted to a neutron plus a high energy positron (β= or e=), increasing the n/p ratio. " " 𝑝 → "6 𝑛 + 6" 𝛽 Process leaves A unchanged, Z decreases by 1. 7Beà7Li + β= + Energy 𝑡"/# = 53 𝑑𝑎𝑦𝑠 144Gd à 144Eu + β= + Energy 𝑡"/# = 4.5 𝑚𝑖𝑛 Trinity College Dublin, The University of Dublin 5. Electron capture Nucleus captures electron from electron cloud " " 𝑝 + 8 6" 𝛽 → "6 𝑛 A is unchanged, Z decreases by 1, as in β+ emission. 51Cr+𝑒8à51V+Energy 𝑡"/# =28𝑑𝑎𝑦𝑠 Trinity College Dublin, The University of Dublin β em ission Positron emission or electron capture α emission NZ = 1 Stable Radioactive à## $%&'()* 23 Nuclear Chemistry But why? unstable isotopes. A graph of the number of neutrons (N) versus the number of protons (Z) for stable (black circles) and radioactive (red All elements beyond bismuth (Z = 83) are unstable. Need to decrease atomic number-> alpha emission FIGURE !"." Stable and 140 circles) isotopes from hydrogen to 110 bismuth. This graph is used to assess 100 90 80 70 60 50 40 30 20 10 00 10 20 30 40 50 60 70 80 90 100 criteria for nuclear stability and to predict modes of decay for unstable nuclei. 130 120 Isotopes above band of stability have high neutron-proton ratio Need to lower neutron-to- proton ratio -> beta emission Isotopes below band of stability have a low neutron-to-proton ratio Need to lower atomic number (move closer to band of stability) Decay by positron emission or by electron capture Number of protons (Z) All elements beyond bismuth (Z = 83) are unstable. To rea stability starting with these elements, a process that decreases t ber is needed. Alpha emission is an effective way to lower Z, the because each emission decreases the atomic number by 2. For cium, the radioactive element used in smoke detectors, decay 243Am → 4α + 239Np 95 2 93 Beta emission occurs for isotopes that have a high neutron-to that is, isotopes above the band of stability. With β decay, the increases by 1, and the mass number remains constant, resu Co→− β+28Ni Number of neutrons (N) Trinity College Dublin, The University of Dublin neutron-to-proton ratio: 60 0 60 Nuclear stability Modes of Radioactive Decay Modes Isotopes Symbol Alpha decay Heavy isotopes 𝛼 or #"𝐻𝑒"& Beta decay Neutron rich isotopes β' or 𝑒' Gamma-ray emission Decay of nuclear excited 𝛾 states Positron emission Proton rich isotopes β& Electron capture Proton rich isotopes X-rays Trinity College Dublin, The University of Dublin Radioactive half-life (𝑡!/#) Time taken for activity to reach half of its initial value 𝑡)/3 = 45 3 𝑁 = 7 = 78!/# 6 69:3 k = reaction rate constant N= number of nuclides A=activity Can relate activity to number of nuclides in sample if t1/2 or k is known. Trinity College Dublin, The University of Dublin Effects of radioactive emissions Nuclear changes cause chemical changes in matter. Virtually all radioactivity causes ionization Number of cation-electron pairs related energy of radiation Resulting charged species go on to form free radicals, molecular or atomic species (with unpaired electrons). Effect on living tissue depends on the penetrating power and ionizing ability of the radiation. Trinity College Dublin, The University of Dublin Units of Radiation Dose (in SI units) Gray (Gy)= 1 joule of energy absorbed per kg ⟹ 1 Gy = 1 J/kg Common unit is the rad (radiation absorbed dose) 1rad=0.01Gy Tissue damage depends on: 1. radiation strength 2. exposure time 3. tissue type 1Gy=1 ( =1 +! )* ,! multiply rads by Relative Biological Effectiveness (RBE) to get radiation dose in roentgen equivalents for man (rem) No. of rems = no. of rads x RBE. (SI unit: 1 Sievert (Sv) = 100 rem) Give us measures of effect of low level ionising radiation on the human body Trinity College Dublin, The University of Dublin s of radiation The S-shaped response model implies there is a threshold above which the effects are more significant: very low risk at low dose and high risk at high dose. Table 24.8 Acute Effects of a Single Dose of Whole-Body Irradiation Dose (rem) Effect 5–20 Possible late effect; possible chromosomal aberrations 20–100 Temporary reduction in white blood cells Lethal Dose 501 Temporary sterility in men (1001 rem 5 1-yr duration) Population (%) ———— — 50–70 60–95 100 No. of Days ———— — 30 30 2 100–200 “Mild radiation sickness”: vomiting, diarrhea, tiredness in a few hours Reduction in infection resistance Possible bone growth retardation in children Permanent sterility in women 3001 500 “Serious radiation sickness”: marrow/intestine destruction 400–1000 Acute illness, early deaths 30001 Acute illness, death in hours to days Trinity College Dublin, The University of Dublin Radiotracers Isotopes exhibit similar chemical and physical behaviour. Certain compounds localize in specific tissues in the body. Map 𝛾 emission from organ of interest e.g. thyroid ##$𝑇𝑐 →##𝑇𝑐+𝛾. I-&TcO-actalike !"!" 4 Relatively safe, no biological function, rapidly excreted Short half-life (6 h), decays to fairly stable nuclei Production of 99mTc pharmaceuticals is very important industry Trinity College Dublin, The University of Dublin 1080 PET Can be used to track a substance in the body – Positron Emission Tomography (brain structure and function) Substance synthesized with isotope that emits positrons. Injected into patient as part of biologically active molecule Positron emitted Patient placed in instrument that measures positron emission. Computer reconstruction of organ Chapter 24 Nuclear Reactions and Their Applications Normal Brain Activity Alzheimer's Disease Figure 24.14 PET and brain activity. These PET scans show brain activity in a normal person (left ) Can be used to measure blood flow, blood volume, oxygen usage, tissue pH (acidity), and in a patient with Alzheimer’s disease (right ). Red and yellow indicate relatively high activity within glucose (sugar) metabolism, and drug activity. a region. Trinity College Dublin, The University of Dublin Additional Applications of Ionizing Radiation Radiation Therapy Cells can be damaged by high energy radiation (Cancer cells more susceptible than health cells) Can be used for treatment cancer inside/outside the body Almost all cases use γ radiation (strongly penetrating) Difficult to avoid healthy cell damage with γ radiation Implants of 198Au or of 90Sr (which decay to a γ emitter) used to destroy pituitary and breast tumour cells, γ rays from 60Co used to destroy brain tumors Effective in slowing cancer or even prompting the regression Trinity College Dublin, The University of Dublin Summary Nuclear decay via three main types of radiation. Half lives and/or decay constants can be calculated. Radiation damages tissue by formation of free radicals. α emitters such as 99mTc very good for imaging. β- emitters such as 90Y useful for cancer therapy. β+ emitters finding more uses in imaging (PET).

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