periodic trends, valence electrons, lewis dot structures, ionic/covalent/metallic bonding

parts of periodic table

• groups: vertical columns, also called families

• periods: horizontal rows

• medium blue (1st column excluding hydrogen): alkali metals

• light purple: alkaline earth metals

• red: transition metals

• orange: metals

• yellow: metalloids

• light blue: nonmetals

• light green: halogens

• darker green: noble gases

• dark blue: lanthanides

• dark purple: actinides

octet rule

• atoms gain or lose electrons (react) to get a full valence shell, which is 8 electrons in most atoms and 2 for hydrogen and helium

• noble gases tend not to react because they already have full valence shells and don’t need any more/fewer electrons

electron configuration down a group

• every atom in a group has the same number of valence electrons

• this leads to a number of shared properties

trend: atomic radius

• size of an atom

• increases down a group

  • adding energy levels → shields outer electrons from attraction to nucleus (opposite charge), creating a weaker nuclear pull

• decreases across a period

  • filling up same energy levels → more electrons added to shell + more protons in the nucleus = larger, more positive nucleus has increased pull on electron cloud, leading to a denser atom

trend: ionization energy

• energy required to take an electron away from an atom

• decreases down a group

  • more energy levels = weaker pull → easier to remove electrons from atoms with more energy levels

• increases across a period

  • extra electrons + bigger, more positive nucleus = stronger nuclear pull → harder to remove electrons

trend: electronegativity

• tendency of an atom to attract electrons in a chemical bond between atoms

• decreases down a group

  • shielding from more energy levels creates weaker nuclear pull → harder to attract electrons

• increases across a period

  • stronger nuclear pull makes it easier to attract electrons

  • think of octet rule: atoms with more valence electrons want to gain rather than lose - e.g. oxygen has six valence electrons so it’s easier to gain two than lose six

trend: electron affinity

• energy released when gaining an electron

• decreases going down a group

  • harder to gain electrons and thus less favorable

• increases across a period

  • stronger pull makes it easier to gain an electron

ionic bonding definition

• intramolecular (within molecules)

• complete transfer of valence electrons

  • usually metals lose electrons, becoming cations

  • usually nonmetals gain electrons, becoming anions

• electronegativity difference must be greater than 2.0

• opposite charge between cat and anions leads to strong attraction - this is the ionic bond

ionic formula

• write metal symbol first

• determine the ionic charge of both

• figure out the number of each using the least common multiple

  • i.e. cation with +2 and anion with +3 needs three cations and 2 anions

type 2 metals

• pretty much overlap with transition metals

• can have more than one ionic charge

  • Fe²⁺ and Fe³⁺ for example

• written with roman numerals in parentheses that represent charge not quantity

  • Iron (II) means iron with 2+ charge not 2 iron atoms

naming ionic compounds

• need to have the formula first

• write name of cation first (use roman numerals if type 2)

  • sometimes suffix ‘ous’ is used for lower charge and ‘ic’ for higher charge for type 2 metals

• add anion name with the ‘ide’ suffix

properties of ionic bonds

• crystalline lattice structure - will shatter easily because of repulsion between like charges when ions get shifted

• high melting and boiling points

• very good electricity conductors but only when molten and aqueous (dissolved in water)

• soluble in water (and dissociates in water)

covalent bonding definition

• intramolecular bonding

• sharing of valence electrons between atoms

  • usually 2 nonmetals

  • between atoms with similar pulls - less than 2.0 electronegativity difference

• in each diagram, a line represents 2 electrons

  • single bond: sharing one pair of electrons (2 e-)

  • double bond: sharing two pairs (4 e-)

  • triple bond: sharing three pairs (6 e-)

formal method for determining formula

• count the number of valence electrons (add or remove based on ionic charge)

• determine the central atom → 1st in the formula name or least electronegative (never hydrogen)

• form single bond between central atom and all others

• distribute remainder of e- equally around the noncentral atoms then dump the rest on the central atom

• check all octets - if needed replace pairs of electrons to form double or triple bonds

formal method addendums (🤪)

• hydrogen only wants 2 e-

• some elements can have more than 8 e- because of empty d sublevel

  • commonly P, S, Si, Cl

• some only need 6 e- instead of a full octet

  • B and Be

• put brackets [ ] around lewis dot structures of ions

naming covalent compounts

• find formula first

• write name of central/least electronegative atom first

  • only put numerical prefixes if there is more than one e.g. water is dihydrogen monoxide but CO@ is just carbon dioxide

• for second atom, use numerical prefix and -ide suffix

properties of covalent molecules

• brittle solids

• usually low melting and boiling points - tend to be liquids and gases

• poor heat and electricity conductors

• ‘like dissolves like’ - polar molecules only soluble in polar substances and likewise for nonpolar

polar vs. non-polar bonds

• non-polar: equal-ish electronegativity (0 - 0.5 difference) leads to equal electron sharing

• polar: less equal electronegativity (0.6 - 1.9) leads to unequal sharing of electrons

  • not fully positive and negative like ionic but one part is partially negative and one part is partially positive

• can calculate bond type by calculating electronegativity difference - the bigger the difference, the more polar the bond

polar vs. non-polar molecule

• IF molecule has non-polar bonds

  • the molecule is always non-polar

• IF molecule has polar bonds

  • AND is symmetrical

    • the molecule is non-polar because molecule is more negative in all directions and thus equally negative

  • AND is asymmetrical

    • the molecule is polar

dipole

• a molecule with two poles (regions with opposite charges)

• dipole arrow represents this, with the cross at the more positive side and pointing to the more negative part

  • -| - - >

  • δ⁺ represents the more positive side and δ^- represents the more negative side

molecular geometry

• electron domain: bonds/bonding pairs (2 e- part of a covalent bond), lone pairs, and unpaired electrons

  • multiple bonds (double/triple) count as one domain

    • multiple bonds also count as one bond in the VSEPR model (valence shell electron pair repulsion theory)

• lone pair electrons take up more space than covalent bonds so lone pairs ‘squash’ bond angles slightly (about 2.5 degrees)

• solid wedge represents bond in front of the page and dashed wedge represents bond behind the page