Valence Electrons

1. Valence Electrons • Electrons in outermost (highest E) shell • Determine chemical properties of an atom • Use to predict how many bonds an atom will make • Group # = number of valence electrons • Can use electron-dot (Lewis-dot) structures to show valence electrons

2. Electron-Dot Structures ofFirst 20 Elements

3. Octet Rule • All noble gases have 8 valence electrons (except He, which has 2) • Noble gases do not combine with other elements because they have a full valence electron shell (very stable) • Octet rule = tendency for atoms to gain, lose or share electrons to obtain a noble gas electron configuration (usually 8 valence electrons) • Elements other than the noble gases thus combine with other elements to form chemical compounds

4. Ions • Ions are atoms, or groups of atoms, with a different number of electrons than protons • Thus, ions have an overall charge, (+) or (-) • An atom can gain electrons to become an anion (-) • Or can lose electrons to become a cation (+) • Ionization energy is E required to remove an electron (takes more E to remove 2nd electron…) • Ionic charges of most ions from main group elements can be predicted from group #

7. Ionic Compounds • Cation + anion = ionic compound • Atoms in ionic compounds are held together by ionic bonds (strong attraction of opposite charges) • There is no sharing of electrons, one atom takes electrons from the other • Charges always balance in ionic compounds (net charge is zero) • It may take more than one atom of either ion to balance the charge • Generally, metals form cations and nonmetals form anions

8. Sodium Gives an Electron to Chlorine to Form Sodium Chloride

9. Magnesium Gives One Electron to Each of Two Chlorines to Form Magnesium Chloride

10. Naming Ionic Compounds • Name of metal cation is just name of metal • Name of anion: replace end of name with ide • Name of cation is followed by name of anion Example: LiBr = lithium bromide • Subscripts are not included in names because they can be assumed (MgCl2 = magnesium chloride) • Some transition metals have variable charges (can’t predict from group #), so amount of charge must be represented in name with a roman numeral: • Fe2+ = iron(II), Fe3+ = iron(III), FeCl2 = iron(II) chloride

11. To Write an Ionic Formula from Name • Write cation from first part of name and anion from second part of name • Balance charges • Write formula with correct subscripts Example: Aluminum sulfide • Al3+ and S2- are the ions • 2 x (3+) = 6+ and 3 x (2-) = 6- (balances charge) • Al2S3 is correct formula

12. Covalent Bonds • Nonmetals form covalent bonds with each other • Covalent bonds involve sharing pairs of electrons • Example: H2 = H:H or H-H (from H. + .H) (two hydrogen atoms form a hydrogen molecule) • Covalent bonds generally follow octet rule (by sharing electrons they gain a full valence shell) • Exceptions to octet rule: - Atoms with 3 valence electrons (B, Al, Ga…) can only make 3 bonds, so only have 6 valence electrons - Some atoms (P, S, Cl, Br, I) can have expanded octets (10, 12 or 14 valence electrons) - H can make only 1 bond, but two electrons gives it a full outer shell (can only have 2 electrons in 1s orbital)

13. Predicting Number of Covalent Bonds from Group Number

14. Multiple Covalent Bonds • Nonmetals can also share 2 or 3 pairs of electrons to form double or triple bonds CO2 is O=C=O and HCN is H-CN • This occurs when necessary to form full octets • C, O, N, S and P are most common atoms to form multiple bonds • H can’t form them, why? (It has only one valence electron to share) • Also, note that H is a nonmetal (even though it’s in Group 1) and does not form ionic compounds as H+)

15. Multiple Bonds in N2 • In nitrogen, octets are achieved by sharing three pairs of electrons • When three pairs of electrons are shared, the multiple bond is called a triple bond octets       N  + N  N:::N  triple bond

16. To Write Dot-Structures from Covalent Formulas • Find central atom (usually forms most bonds) • Find total # of valence electrons (add valence electrons from each atom in compound) • Use a pair of electrons to attach each atom to the central atom • Subtract # of bonded pairs from total and place remaining valence electrons as nonbonded pairs • If not all atoms have a complete octet, move nonbonded pairs to complete (form multiple bonds)

18. Naming Covalent Compounds • More than one covalent compound can often be formed from the same elements (multiple bonds) • Prefixes in the name tell how many of that type of atom are in the compound (subscripts in formula) • Prefixes are optional for single atoms • CO = carbon monoxide (or carbon oxide) • CO2 = carbon dioxide (or monocarbon dioxide) • Naming of organic carbon compounds has a whole unique set of rules (covered in chem 102)

20. Polyatomic Ions • A polyatomic ion is a covalently bonded group of atoms with an overall net charge (number of electrons is different from number of protons) • Most contain covalent bonds between oxygen and another nonmetal (NO3-, HSO3-, PO43-…) • They usually have a charge of -1, -2 or -3 • The only common positively charged one is NH4+

21. Compounds Containing Polyatomic Ions

22. Naming Polyatomic Ions • The ending ate is used for the most common ion for that combination of atoms • Related ions with one less O end in ite PO43- = phosphate PO33- = phosphite • When H is added, add hydrogen to the beginning of the name and reduce negative charge by 1 HPO42- = hydrogen phosphate HPO32- = hydrogen phosphite

24. Writing Formulas with Polyatomic Ions • All ions exist in pairs with ions of opposite charge • Net charge of a compound or a solution is zero • Bonds between polyatomic ions and other ions (mono- or polyatomic) are ionic bonds • Bonds within polyatomic ions are covalent, so the atoms stay together as a group, even when dissolved in water • Use same rules for writing formulas for ionic compounds containing polyatomic ions as those for monoatomic ions (write ions, balance charge, write formula) • Example: magnesium sulfate = Mg2+ + SO42- = MgSO4

25. Shapes of Molecules • Molecules have 3-D shapes • Shape is important to chemical reactivity • Enzymes (proteins) in your cells bind to molecules with specific shapes • Molecular shape can be predicted using VSEPR (valence-shell electron pair repulsion) theory • VSEPR = electron groups (bonds or non-bonded pairs) are placed as far apart as possible around central atom to minimize repulsion between negative charges • Shape is determined by number of electron groups

27. How to Predict Molecular Shape • Write electron-dot structure • Determine number of electron groups and shape 2 groups = linear 3 groups = trigonal planar 4 groups = tetrahedral • If there are lone pairs, remove to determine final shape 3 groups (one lone pair) = bent 4 groups (one lone pair) = pyramidal 4 groups (two lone pairs) = bent

28. Bond Polarity • The electrons in covalent bonds are not always shared equally • Some atoms attract the electrons more • The ability of an atom to attract shared electrons through a covalent bond is called electronegativity • Higher electronegativity = stronger pull on electrons • Fluorine is the most electronegative element (4.0), and values for all other elements are based on F • When electrons in a covalent bond are not equally shared, the bond is called polar covalent

31. Predicting Bond Type from Electronegativity Difference • Bond type is a continuum, from ionic to nonpolar covalent, with polar covalent in the middle • In general: - A metal plus a nonmetal = ionic bond - For a nonmetal plus a nonmetal, an electronegativity difference of: < 0.5 = nonpolar covalent bond ≥ 0.5 = polar covalent bond • Use the symbols + and - to indicate partial charge ( = delta) on atoms in polar covalent bonds • + = electron poor and - = electron rich

34. Polar and Nonpolar Molecules • Molecules can be overall polar or nonpolar • Look first at the shape: 1. If symmetrical, then nonpolar 2. If non-symmetrical, then: - If has polar bonds, then polar - If has non-bonded electon pairs, then polar - If no polar bonds and no lone pairs, then nonpolar

35. A polar molecule contains polar bonds • The separation of positive and negative charge is called a dipole • In a polar molecule, dipoles do not cancel  +  - • • H–Cl H– N–H dipole H dipoles do not cancel

36. A nonpolar molecule contains nonpolar bonds Cl–Cl H–H or a symmetrical arrangement of polar bonds O=C=OCl Cl–C–Cl Cl dipoles cancel