Bonding

1. Bonding Ionic Bonding

2. What is an ionic bond? • A bond formed by the electromagnetic attraction between ions of opposite charge. • Ions can be single charges atoms OR groups of atoms that carry a charge (called polyatomic ions) • Formed between an metal and a nonmetal

3. Oxidation State • A shorthand notation that indicates how an atom is likely to bond. • 2+ indicates that an atom is likely to use or lose 2 electrons when bonding • 1- indicates that an atom is likely to use or gain 1 electron when bonding • Metals have positive oxidation states • Nonmetals have negative oxidation states • Where on the periodic table do you find metals? • Where do you find nonmetals? http://www.chemguide.co.uk/inorganic/redox/oxidnstates.html

4. Where are they located?

6. Criss-Cross to predict chemical formulaIonic Bonding: Magnesium and Iodine (example)

7. Naming Ionic Compounds • Ionic compounds consist of cations (positive ions) and anions (negative ions) held together by electrostatic attraction • Metal + Nonmetal • Ex: Sodium Chloride (salt) = NaCl • When naming ionic compounds, the anion will always end in “-ide” • Ex: Sodium + Chlorine = Sodium Chloride • Ex: Potassium + Iodine = Potassium Iodide

8. Naming monatomic cations • Metal atoms lose valence electrons to form positively charged ions called cations. • An ion formed from an individual atom is a monatomic (or monoatomic) cation.

9. Naming Ionic Compounds That Need A Roman Numeral • The name of an ionic compound must include roman numerals if: • An ionic compound starts with a metal that is not in Group I or Group 2 • This is because many of the metals in the transition metals and the metals under the stairs (the metalloid line) often have multiple oxidation numbers. • For example, iron can take both a "+3" oxidation number and a "+2" oxidation number. • The only way to tell the difference is to use a Roman numeral. We write either Fe(III) or Fe(II) to indicate whether iron has the "+3" or "+2". There are a few exceptions. For example, aluminum is always "+3" and zinc is always "+2". However, in general, you are safer using a Roman numeral if the metal is in this region of the periodic table.

10. Here are the rules: EXAMPLE: Since titanium is a transition metal, it can take more than one possible oxidation number. So we have to do some work to figure out which charge. We know that oxygen usually has a -2 oxidation number: Since there are two oxygen ions, that adds up to -4. Therefore, the Titanium must have a +4 oxidation number to balance the compound. The name is: Titanium (IV) Oxide • Write down the name of the metal • Write down the name of the non-metal • Change the ending of the nonmetal to “-ide” • Use the negative charge on the nonmetal to calculate the oxidation number of the metal • Write the roman numeral that represents the oxidation number of the metal

11. More examples:

12. Interact with Ions • Go to this interactive and follow your notes to guide you to collect more information about ionic bonds! http://kcts9.pbslearningmedia.org/asset/lsps07_int_ionicbonding/

13. Bonding Covalent Bonding

14. Why do atoms bond? • Atoms gain stability when they share electrons and form covalent bonds. • Lower energy states make an atom more stable. • Gaining or losing electrons makes atoms more stable by forming ions with noble-gas electron configurations. • Sharing valence electrons with other atoms also results in noble-gas electron configurations.

15. Why do atoms bond? (cont.) • Atoms in non-ionic compounds share electrons. • The chemical bond that results from sharing electrons is a covalent bond. • A moleculeis formed when two or more atoms bond.

16. Why do atoms bond? (cont.) • Diatomic molecules (H2, F2 for example) exist because two-atom molecules are more stable than single atoms.

17. Why do atoms bond? (cont.) • The most stable arrangement of atoms exists at the point of maximum net attraction, where the atoms bond covalently and form a molecule.

18. Formation of Bonds Atoms are attracted to each other by the opposite charges of their electrons and the other atom’s protons. Although the electrons of the two atoms repel each other, the attraction forces are greater

19. Formation of Bonds At a certain distance, the protons of the two atoms start to repel each other The balance between the attraction and the repulsion results in an ideal bond length between the two atoms.

20. Single Covalent Bonds • When only one pair of electrons is shared, the result is a single covalent bond. • The figure shows two hydrogen atoms forming a hydrogen molecule with a single covalent bond, resulting in an electron configuration like helium.

21. How do we determine the shape in covalent molecules? Covalent molecules are individual units, so they will have specific numbers of atoms instead of general ratios (ionic compounds). That means the empirical formula (lowest ratio of atoms) can be different from the molecular formula (actual number of atoms in a molecule).

22. Lewis Structures—represent a chemical formula, showing unshared and shared valence electrons

23. Single Covalent Bonds (cont.) • In a Lewis structure dots or a line are used to symbolize a single covalent bond. • The halogens—the group 17 elements—have 7 valence electrons and form single covalent bonds with atoms of other non-metals. • Example: HCl, Cl2

24. Single Covalent Bonds (cont.) • Atoms in group 16 can share two electrons and form two covalent bonds. • Water is formed from one oxygen with two hydrogen atoms covalently bonded to it .

25. Single Covalent Bonds (cont.) • Atoms in group 15 form three single covalent bonds, such as in ammonia.

26. Single Covalent Bonds (cont.) • Atoms of group 14 elements form four single covalent bonds, such as in methane.

27. Multiple Covalent Bonds • Double bonds form when two pairs of electrons are shared between two atoms. • Triple bonds form when three pairs of electrons are shared between two atoms.

28. Structural Formulas (cont.) • Drawing Lewis Structures • Predict the location of certain atoms. • Determine the number of electrons available for bonding. • Determine the number of bonding pairs. • Place the bonding pairs. • Determine the number of bonding pairs remaining. • Determine whether the central atom satisfies the octet rule.

29. Structural Formulas (cont.) • Exceptions to the normal method of drawing Lewis structures: • Atoms within a polyatomic ion are covalently bonded. • Ions will change the number of electrons that are used in the Lewis Structure: • Negative ions will have that many more electrons, positive ions will have that many less electrons • Examples: OH-1, NH4+1

30. Exceptions to the Octet Rule • Some molecules do not obey the octet rule. • A small group of molecules might have an odd number of valence electrons. • NO2 has five valence electrons from nitrogen and 12 from oxygen and cannot form an exact number of electron pairs.

31. Exceptions to the Octet Rule (cont.) • A few compounds form stable configurations with less than 8 electrons around the atom—a suboctet. • A coordinate covalent bond forms when one atom donates both of the electrons to be shared with an atom or ion that needs two electrons.

32. Shapes of Molecules That allows for many possible shapes around a central atom: 1. Linear

33. Shapes of Molecules 2. Trigonal Planar

34. Shapes of Molecules 3. Tetrahedral

35. Shapes of Molecules 4. Trigonal Pyramidal

36. Shapes of Molecules 5. Bent

40. Electronegativity and Bond Character • Electronegativity measures the ability of an atom to attract electrons within a bond • There are three types of chemical bonds that can form between atoms: • Ionic bonds • Polar Covalent bonds (partly share electrons) • Non-polar Covalent bonds (fully share electrons)

41. Electronegativity and Bond Character Ionic compounds only form under certain circumstances. When one atom is much more electronegative than another, it will completely take the electron, forming ions (and therefore ionic compounds) When two atoms have similar electronegativity values, they may share the electrons to varying degrees, forming covalent bonds. Atoms covalently share electrons when difference between their attraction is not great (electronegativity diff. less than 1.7)

42. Electronegativity and Bond Character (cont.) • This table lists the character and type of bond that forms with differences in electronegativity. • Noble gases are not listed because they generally do not form compounds.

43. Polar Covalent Bonds • Polar covalent bonds form when atoms pull on electrons in a molecule unequally. • Electrons spend more time around one atom than another resulting in partial charges at the ends of the bond called a dipole.

44. Polar Covalent Bonds (cont.) • Covalently bonded molecules are either polar or non-polar. • Non-polar molecules are not attracted by an electric field. • Polar molecules align with an electric field.

45. Polar Covalent Bonds (cont.) • Compare water and CCl4. • Both bonds are polar, but only water is a polar molecule because of the shape of the molecule.

46. Polar Covalent Bonds (cont.) • Solubility is the property of a substance’s ability to dissolve in another substance. • Polar molecules and ionic substances are usually soluble in polar substances. • Non-polar molecules dissolve only in non-polar substances.

47. Covalent vs. Ionic Bonds Notes Ionic Covalent Contain ions from Electrons are gain or loss of shared (surround electrons both nuclei)

48. Covalent vs. Ionic Bonds Notes Ionic Covalent Each ion has its own Atoms must share complete outer octet electrons to have a complete octet

49. Properties of Ionic Compounds vs. Covalent Compounds Ionic Covalent Strong bonds Weaker bonds

50. Covalent vs. Ionic Bonds Notes Ionic Covalent Form crystals, which are Form individual units a result of strong attraction called molecules in repeated patterns Are hard and brittle Are frequently liquids or gases