Chapter 22 – Chemical Bonds

1. Chapter 22 – Chemical Bonds Chapter preview sections 1 Stability in Bonding 2 Types of Bonds 3 Writing Formulas and Naming Compounds Elements Form Chemical Bonds Just as skydivers link together to make a stable formation, the atoms in elements can link together with chemical bonds to form a compound. You will read about how chemical bond form and learn how to write chemical formulas and equations

2. Chemical Bonds Section 1—Stability in Bonding

3. Combined Elements • Some matter around you is in the form of uncombined elements, such as copper, sulfur and oxygen. • They can unite chemically to form a compound, when conditions are right. • The green coating, on the statue of Liberty is such a compound, called copper sulfate.

4. New Properties • An observation you will make, is that the compound formed when elements combine, often has chemical and physical properties that aren’t anything like those of the individual elements. • Na + Cl  Na+ Cl- Sodium + Chlorine  Sodium Chloride (Table Salt)

5. Formulas • The chemical symbols Na and Cl represent the elements sodium and chlorine. • When written as NaCl, the symbols make up a formula, or chemical shorthand, for the compound sodium chloride. • A chemical formula tells what elements a compound contains and the exact number of atoms of each element in a unit of that compound.

6. Learn Table 1 on page 689

7. The electric forces between electrons and protons, which are opposite, are the forces that cause compounds to form.. The Unique Noble Gases can be understood if you look at an electron dot diagram of them. Atomic Stability

8. An atom is chemically stable when its outermost energy level has the maximum number of electrons. Helium and Hydrogen are stable with two electrons All other elements are stable when they have eight electrons. Stable Unstable Chemical Stability

9. Atomic components • When you look at the elements in Groups 13 through 17, you see that none of the elements has a stable energy level. • Each group contains too few electrons for a stable level of eight electrons. 1 2 13 14 15 16 17 18  Unstable 

10. Hydrogen does not have a full outer energy level. How does it or any other such element, become stable? Atoms with partially stable energy levels can lose, gain, or share electrons. They combine with other atoms that also have partially complete outer energy levels. Ea: Sodium will lose one electron, while chlorine gains one electron, making the combination stable. Outer Levels – Getting Their Fill

11. Stability is Reached • Sodium had only one electron in its outer energy level which it lost when it combined with chlorine to for sodium chloride. • An atom that has lost or gained an electron is called an ion. • An ion is a charged particle because it now has either more or fewer electrons than protons. • It is this electric force that holds these compounds together. • Water (H2O) • In water, Hydrogen and oxygen each contribute one electron to each hydrogen—oxygen bond. The atoms share those electrons instead of giving them up.

12. Types of Bonds Section 2

13. Gain or Loss of Electrons • Atoms can loan electrons to another atom, so they both become stable. • An attractive force occurs with this and we call it a chemical bond.

14. A Bond forms • This type of sharing occurs between atoms in Group 1 and Group 17. The atoms become ions, with one being positive and the other negative. Ea: sodium (Na) joins with chlorine (Cl) to create sodium chloride [Na]+ [Cl]- NaCl

15. The Ionic Bond • When ions attract in this way, a bond is formed. • An ionic bond is the force of attraction between opposite charges of ions in an ionic compound.

16. Zero Net Charge • The result of this bond is a neutral compound. • The compound as a whole is neutral because the sum of the charges is zero. Magnesium + 2 chlorine  Magnesium Chloride

17. Sharing Electrons • Some atoms are unlikely to lose or gain electrons. • Group 14 have 4 electrons in outer shell. • They become more stable by sharing atoms , since removing one makes the atom hold others more tightly.

18. Single Covalent Bond • Bonds formed by sharing are called covalent bonds. • A neutral particle that results is called a molecule. • A single covalent bond is usually made up of two shared electrons. • Water contains two single covalent bonds.

19. Multiple Bonds • A covalent bond also can contain more than one pair of shared electrons. • An example is the bond in nitrogen (N2) • Each atom contributes three electrons. • Each pair of electrons represents a bond.

20. Unequal Sharing • Electrons are not always shared equally between atoms in a covalent bond. • The strength of attraction of each atom to its electrons is related to size of atom. • One example is found in a molecule of hydrogen chloride (HCl). • Chlorine atoms have a stronger attraction than Hydrogen.

21. Partial positive charge δ + Partial negative charge δ - • Tug-of-War: • You might think of a covalent bond as the rope in a tug-of-war, and the shared electrons as the knot in the center. Each atom in the molecule attracts the electrons they share. Some atoms aren’t the same size. Therefore these stronger atoms pull harder in their direction. Hydrogen Fluoride Chloroform

22. Polar or Nonpolar? • For molecules involved in this tug-of-war, there is another consequence. The atom holding the electrons more closely will have a slightly negative charge, and the other atom will be slightly positive. • This type of molecule is called polar. A polar molecule is one that has a slightly positive end, and a slightly negative end. • Water is an example, and its polarity is responsible for many of its unique properties

23. Polar or Nonpolar? • Two atoms that are exactly alike can share their electrons equally, forming a nonpolar molecule. • This is true of molecules with identical atoms, or symmetric atoms (CCl4) carbon tetrachloride Nonpolar (symmetric) Hydrogen molecule  Nonpolar  Oxygen molecule

24. Properties of Compounds • Recall: Atoms can form two types of bonds, covalent and ionic. • Sugar is a covalent compound • Table salt is an ionic compound • Their comparison can be seen in the table on page 701 of you books.

26. Covalent & Ionic Properties • The difference in properties is due to differences in attractive forces of the bonds. • Covalent Compounds: • Covalent bonds between atoms are strong • Attraction between molecules is weak • Melting & boiling points are relatively low (sugar melts at 185ᵒC. • They form soft solids, which have poor electrical and thermal conductivity. (Ea: candles & propane gas)

27. Covalent & Ionic Properties (continued) • Ionic Compounds: • Bonds between ions are relatively strong. • They have high melting & boiling points – Salt melts at 801ᵒC • Solids are hard and brittle. • When in liquid or aqueous state they readily conduct electric current. • Ionic compounds are stable due to the strength of attraction.

28. Writing Formulas and Naming CompoundsSection 3 The numbers in columns 1, 2, 13, 14, 15, 16, 17 & 18; is the most common oxidation number if elements in that group (family).

29. A binary compound is one that is composed of two elements. Before you can write a formula, you must have all needed information. What you will need to know: Identify the symbols of the cation (first part of the name) and anion Identify the charge for each and place above the symbol in parenthesis Balance the positive and negative charges Write the formula placing the subscripts right after the symbol they go with. Binary Ionic Compounds

30. Oxidation Numbers • You will need to know which elements are involved and what number of electrons they lose, gain, or share to become stable. • Because all elements in a group have the same number of electrons in their outer energy levels, they must gain or lose the same number of electrons. • Metals always lose electrons and nonmetals always gain electrons when they form ions. • The charge on the ion is known as the oxidation number of the atom.

31. Oxidation Numbers & Periodic Table • The numbers with positive or negative signs are the oxidation numbers for the elements in the column below them.

32. Writing Formulas • When writing an ionic formula between a metal and a nonmetal follow these 5 steps: • Write the symbols for the metal and the nonmetal. • Write the valences as superscripts above each symbol. • Drop the + and - sign. • Crisscross the valences so they become the subscript for the other element. • Reduce subscripts whenever possible. Only when both are divisible by a number greater than one. • Let's use these rules to figure out the chemical formula for our compound between aluminum and oxygen.

33. Tables for Writing Names

34. Writing Names (See rules on p.706) • You can name a binary ionic compound from its formula by using these rules: • Write the name of the positive ion. • Use Table 3 – Check to see if there could possibly be more than one oxidation number. • Write the root name of the negative ion. (the first part of the electrons name. Ea: Chlorine becomes chlor. • Add the ending –ide to the root. Table 4 list several elements and their –ide counterparts.

35. Compounds with Polyatomic Ions • Not all ionic compounds are binary. • Baking soda has the formula: NaHCO3. • This is an example of a non-binary ionic compound. • A polyatomic ion is a positive or negatively charged, covalently bonded group of elements. • The polyatomic ion in Baking soda is the bicarbonate or hydrogen carbonate ion HCO3.

36. Polyatomic ions (continued) • Writing Names • Several Polyatomic ions are named in Table 5. • First write the name of the positive ion. • Then use table 5 to find the name of the polyatomic ion. • Ea: K2SO4 is Potassium sulfate, • Sr(OH)2 is Strontium hydroxide

37. Polyatomic ions (continued) • Writing formulas • Follow the rules for binary compounds, with one addition. When more than one polyatomic ion is needed, write parentheses around the ion before adding the subscript. • Barium chlorate: Barium is Ba2+ and chlorate is ClO31- • The formula is: Ba(ClO3)2

38. Compounds with Added Water • Some ionic compounds have water molecules as part of their structure. These compounds are called hydrates • A hydrate is a compound that has water attached chemically to its ions and written into its chemical formula. • Common Hydrates: Copper(II) sulfate pentahydrate, • CuSO4· 5H2O. The name and formula indicate that there are five water molecules per copper sulfate formula unit. • Hydrated water molecules are generally indicated in formulas as shown using a dot to separate the water molecules from the formula of the salt. • If copper(II) sulfate pentahydrate is heated, the bright blue crystals of the hydrate are converted to a white, powdery, anhydrous salt. Anhydrous – “without” water

39. Naming Binary Covalent Compounds • Covalent compounds are those formed between elements that are nonmetals. • For example: Nitrogen and oxygen can form N2O, NO, NO2, N2O5. • These would all be called nitrogen oxide. • Using Prefixes: Using Greek prefixes, these would be called dinitrogen oxide, nitrogen oxide, nitrogen dioxide, dinitrogen pentoxide.

40. Naming Binary Covalent Compounds(continued) • Notice that the last vowel of the prefix is dropped when the second element begins with a vowel. • Carbon monoxide is an example. • These same prefixes are used when naming the hydrates previously discussed.