Periodic Trends and Atomic Properties

• Course FSC112/CHM 101 topics include trends in atomic radii, ionization energies, electronegativity, chemical equations, stoichiometry, calculations based on equations, and the limiting reactant concept.

Periodic Table

• The layout is a tabular arrangement of elements by increasing atomic number.
• Periodicity refers to the recurrence of similar properties at regular intervals with elements that have increasing atomic numbers.
• Main group and transition elements are sometimes labeled as A and B subgroups.
• s-block and p-block groups numbered as I to VII as well as zero.
• d-block and f-block elements also exist.
• 19th-century chemists like Newland, Lother, and Mendeleev contributed to the modern periodic table.
• Elements are arranged by increasing atomic number.
• Elements in the same group (vertical column) have similar properties and the same number of valence shell electrons.
• Lithium (Li) has the electron configuration [He] 2s¹.
• Sodium (Na) has the electron configuration [Ne] 3s¹.
• Potassium (K) has the electron configuration [Ar] 4s¹.
• Rubidium (Rb) has the electron configuration [Kr] 5s¹.
• Elements in the same period (horizontal row) have progressively different physical and chemical properties as electrons and protons are added.
• Lithium (Li): 3, [He]2s¹
• Beryllium (Be): 4, [He]2s²
• Boron (B): 5, [He] 2s²2p¹
• Carbon (C): 6, [He]2s²2p²
• Ground state electron configurations describe the electron arrangement outermost energy level.

Periodicity

• Periodicity involves recurring elemental properties due to regular, predictable variations in atomic structure.
• Periodic properties depend on electronic configuration and repeat at certain intervals within groups or periods.
• Gradual changes in physical properties are shown by elements across a period or down a group.
• Properties include atomic radius, ionic radius, ionization energy, electron affinity, and electronegativity.

Atomic Radius

• Half the distance between the nuclei of two atoms in a homodiatomic molecule is the atomic radius.
• Atomic radius decreases across a period due to increased positive charge from more protons in the nucleus.
• A greater positive charge pulls added electrons closer to the nucleus.

Trends in Atomic Radii

• Energy level of the valence electron, nuclear charge, and shielding effect which influence it.
• Higher energy levels result in valence electrons that are further from the nucleus.
• Distance increases by energy level: 1 < 2 < 3 < 4 < 5 < 6 < 7
• Increased nuclear charge pulls electrons closer to the nucleus.
• Nuclear charge equals atomic number.
• A higher atomic number reduces size.
• Shielding by inner electrons screens outer electrons from the nucleus.
• Greater shielding leads to larger atomic size due to electron repulsion.

Ionization Energy (IE)

• The energy (kJ/mole) needed to remove the most loosely bound electron from an isolated gaseous atom to form a positive gaseous ion.
• Supplying a small amount of energy promotes electrons to a higher energy level.
• A sufficient amount of energy completely removes the electron.
• First ionization energy is the minimum energy needed to remove the first electron from a gaseous atom in its ground state, forming a gaseous ion.
• Subsequent ionization energies(I₂, I₃, etc.) refer to the energy to remove the second, third, and so on electrons.
• Each subsequent IE is progressively greater (I₁ < I₂ < I₃).
• Ionization energy is always endothermic, or requires energy to be added to the atom.
• IE reduces if the size of the atom increases.
• Ionization energy and atomic radius have an inverse relationship.
• Larger atoms have smaller ionization energies and larger atomic radii.
• Nuclear charge, shielding effect, radius and sublevel impact IE.
• Larger nuclear charge increases ionization energy.
• Greater shielding effect lessens ionization energy.
• Larger distance between nucleus and outer electrons lowers ionization energy.
• Removing an electron from a full or half-full sublevel requires extra energy.
• Full and half full orbitals are more stable.
• Filled ns² and np⁶ orbitals along with half-filled np³ orbitals are more stable.
• Higher ionization energy than expected is seen because the orbital resists losing stability
• Group trends are based on first, second, and third ionization energies.
• The first IE decreases down a group as atomic radius increases and shielding effect increases.
• In the same period, atoms have valence electrons in the same energy level and shielding.
• Nuclear charge increases across the period.
• IE generally increases from left to right but exceptions occur with full/half full orbitals.
• In the second and third rows, successive ionization energies rise as electrons are removed from valence orbitals (2s or 2p & 3s or 3p).
• A sharp increase is seen in ionization energy occurs when electrons are taken from filled core levels.
• First ionization energies increase across the periodic table as valence electrons do not screen each other, and increase the effective nuclear charge across the row.
• Valence electrons are more strongly attracted to the nucleus and this reduces atomic size and increases ionization energy.
• Changes seen in the second, third, fourth, fifth, and sixth row, s and p block ionization energies increase across each row from left to right.
• First ionization energies fall down a group.
• Filled inner shells effectively screen valence electrons.
• The small increase is due to an effective nuclear charge.
• Atoms become larger as they gain electrons.
• Valence electrons further from the nucleus are less tightly bound, lessening ionization energies.
• A larger atomic radius has a lower ionization energy.
• Elements forming positive ions most easily are in the lower-left corner having the lowest ionization energies.
• The elements that resist ionization are in the upper-right corner.
• Ionization energies rise diagonally from the lower-left to upper-right.
• Minor deviations are due to stable electronic configurations called pseudo-inert gas configuration.

Electronegativity

• Electronegativity measures an atom's ability to attract electrons within a molecule.
• Linus Pauling (1901-1994) first proposed the concept and later won a Nobel Prize.
• Metals have a low tendency to attract electrons and a high tendency to release them.
• Non-metals have a high tendency to attract electrons and a low tendency to release them.
• Smaller atoms attract electrons more strongly and are more electronegative.
• Atoms with nearly filled electron shells have higher electronegativity that those with sparsely filled shells.
• A covalent bond predominates between two atoms with similar electronegativities.
• A large electronegativity difference leads to a bond with high polar character.
• In a group, larger atoms increase in size, and valence electrons move further from the nucleus.
• This increases shielding from nuclear charge.
• The lessened attraction of the nucleus for valence electrons reduces electronegativity down the group.
• The sizes of atoms decrease across a period, retaining the same energy levels.
• Attraction of the nucleus for valence electrons increases.
• Electronegativity rises left to right across the period
• Fluorine (F) is the highest, most electronegative element.

Chemical Equations

• Shorthand descriptions the chemical reactions
• Symbols and formulas represent the elements and compounds, for example: CH₄ + 2O₂ → CO₂ + 2H₂O.
• Reactants CH₄ and O₂ are on the left side of the arrow, and products CO₂ and 2H₂O are on the right.
• Atoms are rearranged, bonds have been broken, and new bonds have been formed.
• An equation must be balanced from the same number of each type of atom on both the product and reactant side.
• Balancing chemical equations ensures the concept is upheld by ensuring there is an equal number of atoms of each element on both sides of the equation.
• Balances change coefficients, not subscripts to preserve the identities of reactants and products.
• Balancing is usually done by inspecting the equation, and is a trial and error process.
• Complex molecules with the molecule with the greatest number of atoms are started with.

Stoichiometric Calculations

• The equation must be balanced before performing any stoichiometric calculation.
• For the equation, C₃H₈(g) + 5O₂(g) → 3CO₂(g) + 4H₂O(g)
• 1 mole of C₃H₈ reacts with 5 moles of O₂ to produce 3 moles of CO₂ and 4 moles of H₂O

Limiting Reagent

• It is completely consumed during a chemical reaction. The amount of product formed is limited by it.
• The stoichiometry of reactants must be known.
• For example, equation 2Bd + 1 egg → Bd₂egg -Toast bread requires 2 slices of bread (Bd) and 1 egg.
• If the reactant has 20 slices of bread and 12 eggs, the bread is limiting

Percentage Yield

• Mass of product formed in a chemical reaction
• Percent yield is the actual yield ratio to the theoretical yield
• The amount of a formed product when the limiting reactant yield is completely consumed, is called the theoretical yield.
• The actual yield is what you actually get
• Cases where actual yields are not the same a theoretical
• Yield is determined by calculating the percentage. Actual yield divided by the theoretical yield times 100%. When the actual equals theoretical, the reaction process is a stoichiometric process.