Nomenclature- Binary Ionic Compounds

1. Nomenclature- Binary Ionic Compounds Binary compounds are compounds with two different kinds of atoms. Naming Binary Ionic Compounds: name = metal name + stem of nonmetal name + -ide The stem names and ionic symbols for some common nonmetals are given in the following table:

2. Examples of Binary Ionic Compounds Name K2O: name = metal name + nonmetal stem + -ide name = potassium + ox + -ide = potassium oxide Name Mg3N2: name = metal name + nonmetal stem + -ide name = magnesium + nitr + -ide = magnesium nitride Name BeS: name = metal name + nonmetal stem + -ide name = beryllium + sulf + -ide = beryllium sulfide Name AlBr3: name = metal name + nonmetal stem + -ide name = aluminum + brom + -ide name = aluminum bromide

3. Metals forming ions more than one charge Some metal atoms, especially those of transition and inner-transition elements form more than one type of charged ion. Cobalt, for example, forms both Co2+ and Co3+ ions. The binary compounds containing such ions are named following the pattern given earlier with one addition, the number of positive charges on the metal ion is indicated by a Roman numeral in parentheses following the metal name. For example, the compounds CoO and Co2O3 contain cobalt ions with 2+ and 3+ charges respectively. Their names are cobalt (II) oxide and cobalt (III) oxide. An older system for metals that form only different ions- endings �ous and �ic are attached to the stem of the metal name. (Non-English names for elements with symbols derived from non-English names)

4. Nomenclature- Covalent Compounds Covalent compounds most often between representative elements classified as non-metals. The pattern used to name binary covalent compounds is similar to that used to name binary ionic compounds: name = name of least electronegative element + stem of more electronegative element + -ide.

5. Examples of binary covalent compounds SO2: name = sulfur + di- + ox + -ide = sulfur dioxide XeF6: name = xenon = hexa- + fluor + -ide = xenon hexafluoride H2O: name = di- + hydrogen + mono- + ox + -ide = dihydrogen monoxide (also known as water). Note, the final o of mono- was dropped for ease of pronunciation.

6. Ionic Compounds with polyatomic ions The rules for writing formulas for ionic compounds containing polyatomic ions are essentially the same as those used for writing formulas for binary ionic compounds. The symbol for the metal is written first, followed by the formula for the negative polyatomic ion. Equal numbers of positive and negative charges must be represented by the formula.

7. Ionic Compounds with polyatomic ions Examples: Compound containing K+ and ClO3-: Compound containing Ca2+ and ClO3-: Compound containing Ca2+ and PO43-: The names of ionic compounds that contain a polyatomic ion are obtained using the following pattern: name = name of metal + name of polyatomic ion Examples: KClO3 is named potassium chlorate ; Ca(ClO3)2 is named calcium chlorate; Ca3(PO4)2 is named calcium phosphate; CaHPO4 is named calcium hydrogen phosphate.

8. Reading Chapter 3 Revisit Sections 3.5-3.6 Chapter 4 Revisit sections 4.4, 4.10 Read sections 4.1-4.3, 4.6

9. Noble gas configuration Noble gases have completely filled s and p subshells of their valence (outermost) shell (except He) Octet rule - rule for predicting electron behavior in reacting atoms. �Atoms will gain or lose electrons to achieve an outer electron arrangement identical to that of a noble gas.� Learning check 4.1

10. Lewis structure Lewis structure (also called electron-dot formula)- Nucleus and all electrons around the nucleus except those in the valence shell- Elemental symbol Valence shell electrons- Dots arranged around the symbol. Number of valence electrons Valence electrons as those having the largest n value in the configuration. For representative elements- Number of valence electrons same as the Roman numeral group number. Example 4.2, Learning check 4.2

11. Ionic bonding Simple Ions- Atoms acquire a net positive or negative charge by losing or gaining one or more electrons The attractive force that holds together ions of opposite charge is called as ionic bond. Metals lose electrons Non-metals gain electrons Example 4.4 Magnesium, Mg, has two valence electrons which it loses to form a simple ion with a +2 electrical charge. The ion is written as Mg2+. Oxygen, O, has six valence electrons. It tends to gain two electrons to form a simple ion with a -2 electrical charge. The ion is written as O2-. Bromine, Br, has seven valence electrons. It tends to gain one electron to form a simple ion with a -1 electrical charge. The ion is written as Br -.

12. Ionic bonding Representative metals form +vely charged ions:Charge = number (Roman numeral) of the group. Representative nonmetals form �vely charged ions:Charge = 8 - the number (Roman numeral) of the group. Example 4.5, Learning check 4.5 Atoms are changed into ions with noble gas configurations* Ionic bonds- attractive force between oppositely charged ions Ionic Compounds- The substances formed when ionic bonds form between positive and negative ions are called ionic compounds. Formulas- ratio of the +ve and �ve ions ? total charge = zero Example 4.6

13. Covalent Bonding Covalent bonding- sharing of electrons to satisfy octet rule of each atom Example- F2 Representation: shared pair of dots or single line Bonding due to overlap of atomic orbitals

14. Covalent bonding Covalent bond- Attractive force between two atoms both attracted to a shared pair of electrons Sharing between identical atoms (homoatomic): Cl2, O2 and N2 Sharing between different atoms (heteroatomic): H2O, and CH4 Example 4.10, Learning check 4.10, O + O ? O2 (more than two electrons shared)

15. Examples of covalent bonding

16. Lewis structures for covalent molecules Step 1: Molecular formula and number of atoms CO2 (1 atom C, 2 atoms O) Step 2: Initial structure using the connecting pattern O C O Step 3: Total number of valence shell electrons (Use periodic table group number) O=6, C=4, O=6. Total= 16 Step 4: Place a pair of electrons between each bonded pair of atoms. Subtract from total valence electrons. Use remaining to complete octet of atoms O:C:O ? 16-2-2 = 12 electrons left

17. Lewis structures for covalent molecules Step 5: Check octet rule for all atoms. If yes, Lewis structure is complete. If not, move unbonded pairs to positions between bonded atoms to complete octets. Octet of O atoms not complete Moving one unbonded pair of electrons from C atom between the O and C atoms or Double bonds- Sharing of two pairs of electrons Triple bonds- Sharing of three pairs of electrons Example 4.11, Learning check 4.11

18. Reading Revisit- Sections 4.1-4.3, 4.5 Read 4.7-4.9 Quiz 3 (Chapter 3) on Friday, Feb 2nd. HW 3 will be posted tonight (blackboard). Due on Mon, Feb 5th. Exam 1 on Wed, Feb 7th. Chapters 1-4.

19. Shapes of molecules and polyatomic ions not flat 2-D distinctive 3-D shapes Valence-shell electron-pair theory or VSEPR theory Mutual repulsion of electron pairs in valence shells Applied to central atom (bonded to two or more other atoms) of a molecule or an ion to predict the shape Step 1: Draw the Lewis structure Rules for VSEPR: Rule 1: All valence-shell electron-pairs considered (bonding and non-bonding) Rule 2: Double or triple bonds with central atom treated as a single pair for prediction

20. VSEPR theory Two pairs of electrons: One pair on each opposite side of central atom Example: CO2 Central atom: carbon Rule 1: Consider the two electron-pairs each side Rule 2: Treat each double bond as single pair Shape: Two pairs located on opposite sides of C atom. A linear molecule

21. VSEPR theory Three pairs of electrons: Triangle Example: ozone (O3) Central atom: oxygen Rule 1: Consider the four electron-pairs (3 bonding, 1 non-bonding) Rule 2: Treat the double bond as a single pair Shape: The double-bonded pairs, 1 bonding pair and 1 non-bonding pair arrange as a flat triangle

22. VSEPR theory Four pairs of electrons: Tetrahedral (four corners of a tetrahedron) Example: water (H2O) Central atom: oxygen Rule 1: Consider the four electron-pairs (2 bonding, 2 non-bonding) Rule 2: No double bond Shape: The two bonding pairs and 2 non-bonding pairs arrange as four corners of tetrahedron

23. VSEPR theory Five electron-pairs Six electron-pairs Learning check 4.13 Example 4.14

24. Polarity of molecules Homoatomic molecules- shared electron pair attracted equally. Nonpolar covalent bonds Electronegativity is measure of tendency of an atom to attract shared electrons of a covalent bond Results in shifting of shared electrons towards the more electronegative atom called Bond polarization Polar covalent bonds More electronegative- partial negative charge (d-) Less electronegative- partial positive charge (d+)

25. Polarity of molecules Example 4.15, learning check 4.15 Extent of bond polarization- electronegativity differences (?EN) (?EN)=0.0 ? Nonpolar covalent (?EN) =2.1 ? Ionic 0.0<(?EN) <2.1 ? Polar covalent Example 4.16 Learning check 4.16