• ionic bond – electrostatic forces between ions of opposite charges
• transfers electrons
• metals and nonmetals
• covalent bond – sharing of electrons
• nonmetallic elements
• metallic bonding – bonding between metal atoms

• valence electrons involved in chemical bonding
• electron-dot symbol – shows valence electrons of atoms at bonds with other atoms
• octet rule – atoms tend to lose or gain electrons until they have eight valence electrons

• metal of low ionization energy and nonmetal with high affinity
• exothermic
• transfer of electrons

• ionic compounds stable because of attraction between opposite charges of ions
• lattice energy – energy required to separate one mole of a solid ionic compound into its gaseous ions
• ionic compounds are hard, brittle, high melting points
• potential energy of two interacting charged particles:
• (Q₁ and Q₂ = charges of particles, d = distance between centers, k = constant = 8.99 x 10⁹ J-m/C²
• attractive interaction increases as magnitudes of charges increase as distance between centers decreases
• lattice energy increases as charges on ions increase and as radii decrease
• magnitude of lattice energy depends on ionic charges

• never find ionic compounds that contain Na²⁺ and others like it
• increase in lattice energy not enough to compensate for second ionization energy
• addition of electrons to a higher shell is energetically unfavorable

• in forming ions, transition metals lose valence-shell s electrons first, then as many d electrons as are required to reach the charge of the ion

• a polytomic ion acts as one charged specie in forming ionic compounds

• size of ion depends on nuclear charge, number of electrons, outer-shell orbitals
• cations are smaller than parent atoms
• anions larger than parent atom
• ion of same charge, size increases down a group
• as principal quantum number increases, size of ion and parent atom increases
• isoelectronic series – ions having the same number of electrons
• increase in nuclear charge, decrease in atomic radius

• lewis structures: shared electrons shown as lines and unshared as dots

• distance between bonded atoms decreases with increasing shared electron pairs

• bond polarity – describes sharing of electrons
• nonpolar covalent bond – electrons shared equally between atoms
• polar covalent bond – one atoms has greater attraction toward for bonding electrons
• if large ionic bond occurs

• electronegativity – ability of an atom in a molecule to attract electrons to itself
• Linus Pauling – first person to develop electronegativity scale
• Electronegativity increases across period
• Decreases down group

• use difference of electronegativity to determine polarity
• greater difference of electronegativity more polar the bond

• 1) sum the valence electrons from all atoms
• 2) write symbols for atoms and connect with single bond
• 3) complet octets of atoms bonded to central atom
• place leftover electrons on central atom even if more than an octet
• if not enough for octet use double bond or triple bond

• formal charge – charge an atom in a molecule would have if all atoms had same electronegativity
• calculate formal charge:
• 1) all of unshared electrons assigned to original atom
• 2) half bonding electrons assigned to each atom in bond
• formal charge of atom equal number of valence electrons in isolated atom, minus number of electrons assigned to atom in Lewis structure

• alternate Lewis structures of same molecule

• lewis structures of benzene

• 1) molecules with odd number of electrons

- 2) molecules in which an atom has less than an octet

• 3) molecules in which an atom has more than on octet

• ClO₂, NO, NO₂
• Octet cannot be achieved

• compounds of boron and beryllium

• larger cental atom, the larger the number of atoms that can surround it
• expanded valence shells occur when central atom bonded to smallest most electronegative atoms

• bond enthalpy – enthalpy change for breaking a particular bond in a mole of gaseous substance
• bond enthalpy always positive

• 
• bond enthalpies derived for gaseous molecules and are averaged values

• bond length – distance between the nuclei of bonded atoms
• as number bonds increases, the bond grows shorter

• 1) oxidation number of element in elemental from is zero
• 2) oxidation number of monoatomic ion is same as charge
• 3) in binary compounds element with greater electronegativity assigned a negative oxidation umber equal to charge in simple ionic compound of element
• 4) sum of the oxidation numbers equal zero for electrically neutral compound and equals overall charge for ionic species

• naming binary compounds: one for ionic compounds and other for molecular compounds
• less electronegative element is given first, then more electronegative atom with –ide ending
• compounds of metals in higher oxidation states molecular rather than ionic