Chemical bonds

1. Chemical bonds An introduction to chemistry

2. Compounds and chemical change • Atom - smallest elemental unit • Molecule • smallest particle still retaining the characteristic chemical properties of a substance • Examples: • oxygen, hydrogen gas - diatomic molecules • Ozone - triatomic oxygen molecule • Noble gases: helium, neon - “monatomic” molecules

3. Chemical reactions • Formation and/or breaking of chemical bonds to form new molecules (products) from old ones (reactants) • Chemical energy - internal bonding potential energy • Chemical equation - symbolic summary of chemical reaction

4. Magnesium is an alkaline earth metal that burns brightly in air, releasing heat \and light. As chemical energy is released, a new chemical substance is formed. The new chemical material is magnesium oxide, a soft powdery material that forms an alkaline solution in water (called milk of magnesia).

5. A chemical equation is a shorthand representation of a chemical reaction. • The substances that are changed in to reaction are called the reactants. • The substances that are formed in the reaction are called the products. • chemical reactions and energy flow can be explained by the making or breaking of chemical bonds. -- chemical bonds can be explained in terms of changes in the electronstructure of atoms

6. Valence electrons are those electrons that occur on the outer shell of atoms. • These are the electrons that are responsible for chemical reactions • Inner electrons are in lower energy levels and the orbitals are filled. • These electrons are therefore not available for interactions with other electrons.

7. The Octet Rule • Atoms attempt to acquire an outer orbital with eight electrons through chemical reactions (octet rule) • This gives them an outer shell configuration like their nearest noble gas and therefore they become stable. • From the family number of the representative elements, you can determine the number of valence electrons, and therefore the number of electrons necessary to gain the stable configuration • Sodium is in Group IA and therefore has 1 electron in its outer shell. • If sodium loses one electron it becomes Na+ and has 8 electrons in its outer shell. • It then has the electron configuration of Ne (neon) and has a filled outer shell configuration.

8. Valence electrons and ions • Outer electrons determining the chemical properties of an atom • Octet rule • Atoms attempt to acquire an outer shell of eight electrons • Electrons can be gained/lost/shared in the process • Example: sodium (Na)

9. Attractive forces holding atoms together in compounds Can be described in terms of molecular (delocalized) or atomic (localized) orbitals Three types: Ionic Electrons transferred between atoms Electrostatic force = binding force Covalent Octets achieved through sharing electrons Typically between nonmetallic elements, r.h.s of periodic table Metallic bonds Outer electrons move freely throughout metal “Electron gas” within rigid lattice of metal atoms Conduct heat and electricity well Chemical bonds

10. Ionic bonds • Chemical bond of electrostatic attraction • Form crystalline solids with orderly geometric structure • Example: NaCl • Na loses; Cl gains • No single NaCl molecule, per se

11. Sodium chloride crystals are composed of sodium and chlorine ions held together by electrostatic attraction. Each sodium ion is surrounded by six chlorine ions, and each chlorine ion is surrounded by six sodium ions. A crystal builds up like this, giving the sodium chloride crystal a cubic structure.

12. You can clearly see the cubic structure of these ordinary table salt crystals because they have been magnified about ten times.

13. Energy and electrons in ionic bonding • Reaction energy released = heat of formation • Divided conceptually into half-reactions Electron transfer rules • Electrons lost/gained to form closed octets • Number gained = number lost

14. Two rules for keeping track of electrons in ionic bonding reactions. • Ions are formed when atoms gain or lose electrons to achieve a noble gas configuration • The number of electrons that are lost must equal the number of electrons that are gained. • Electrons are not created or destroyed in a chemical reaction

15. Ionic compounds and formulas Formulas • List elements in compound and their proportions • Proportions decided by electron gain/loss Ionic compounds • Characterized by ionic bonds • White, crystalline solids soluble in water • Families IA and IIA lose electrons and form positive ions • Families VIA and VIIA gain electrons to form negative ions

16. Covalent bonds • Chemical bonds formed by sharing pairs of electrons • Electrons shared to form octets, ideally • Overlap of shared electron clouds between nuclei yields net attraction • Atoms within covalent compounds are electrically neutral, or nearly so

17. Covalent Bonds • A covalent bond is a chemical bond that is formed when two atoms share a pair of electrons. • H. + H. H:H • Covalent Compounds and Formulas • Since a pair of electrons is shared in a covalent bond, the electrons move throughout the entire molecular orbital. • In the above example, since both hydrogen share the electron pair, each hydrogen has a filled valence shell, since it has the electron configuration of helium. • Compounds that are held together by covalent bonds are called covalent compounds. • Covalent compounds form from atoms on the right side of the periodic table • Families IVA through VIIA, the number of unpaired electrons (number of covalent bonds) is eight minus family number.

18. Covalent compounds and formulas • Covalent compound - held together by covalent bonds • Electrons shared in covalent bonds • Electron dot representation • Bonding pairs shared • Lone (non-bonding) pairs not shared

19. Multiple bonds • Sharing of more than one electron pair • Examples • Ethylene - double bond • Acetylene - triple bond

20. Acetylene is a hydrocarbon consisting of two carbon atoms and two hydrogen atoms held together by a triple covalent bond between the two carbon atoms. When mixed with oxygen gas (the tank to the right), the resulting flame is hot enough to cut through most metals.

21. Coordinate Covalent Bonds • A coordinate covalent bond is one in which the electron pair comes from one atom. H+ + NH3 NH4+ H H .. .. H+ + :N:H H:N:H .. .. H H • In this example both electrons from the new covalent bond come from the lone pair (non-bonding electrons) around the nitrogen.

22. Bond polarity • Result of unequal sharing of electrons • Electronegativity • Measure of an atom’s ability to attract electrons • Differences: • 1.7 or greater - ionic • 0.5-1.7 - polar covalent • Less than 0.5 - covalent

23. Electronegativities

24. Composition of compounds • Millions of different combinations of over 90 elements • Common names • Often related to historical usage (baking soda, washing soda,…) • Difficult to relate to actual molecular composition • Modern approach - systematic sets of rules • Different for ionic and covalent compounds • One common rule - “-ide” means compound contains only two different elements

25. These substances are made up of sodium and some form of a carbonate ion. All have common names with the term "soda" for this reason. Soda water (or "soda pop") was first made by reacting soda (sodium carbonate) with an acid, so it was called "soda water."

26. Ionic compound names • Name of metal (positive) ion first; then nonmetal (negative) ion • Many elements have variable charges • Historical suffix usage • “-ic” for higher of two; • “-ous” for lower • Modern approach • English name of metal followed by Roman numeral indicating charge

27. Ionic Compound Names • Ionic compounds that are formed from metal ions are named by naming the metal ion (electropositive ion) first, followed by the nonmetal (electronegative ion) • The ending of the nonmetal is changed to end in -ide • When a metal can have various oxidation states the oxidation state is give by roman numerals in parenthesis after the name of the metal. Older naming added –ic & -ous to the end of the name of the metal, for the higher and lower of two possible charges. -- some ionic compounds have polyatomic ions, thereby containing 3 or more elements. ( Metal hydroxides) (Salts: -ate, -ite, per, hypo)

28. Ionic compound formulas • Two rules • Write symbol for positive ion first followed by negative ion symbol • Assign subscripts to assure compound is electrically neutral • Example: Calcium chloride

29. A battery and bulb will tell you if a solution contains ions.

30. Molecular - composed of two or more nonmetals Same elements can combine to form a number of different compounds Two rules First element in formula named first with number indicated by Greek prefix Stem name of second element next; Greek prefix for number; ending in “-ide” for two elements) Covalent compound names

31. Covalent Compound Names • Names for covalent compounds uses Greek prefixes to indicate numbers of atoms of each element • The first element in the formula is named first with a prefix indicating the number of atoms if the number is greater than one. • The stem name of the second element in the formula is named next, with a prefix used with the stem name if two elements can form more than one compound • The suffix –ide is again used.( either ionic or covalent) • Example • CO = carbon monoxide • CO2 = carbon dioxide

32. Covalent Compound Formulas • The systematic name tells you the formula • Formulas indicate how many atoms of one element combine with atoms of another element. • The number of covalent bonds that an atom can form is called its valence. • Lone pairs create the possibility of creating coordinate covalent compounds.

33. Once you understand chemical names and formulas, you can figure out what chemical compounds are contained in different household products. For example, (A) washing soda is sodium carbonate (NA2CO3) and (B) oven cleaner is sodium hydroxide (NaOH), which is also known as lye.

34. Covalent compound formulas • Examples: carbon dioxide, carbon tetrachloride • Valence • Number of covalent bonds an atom can form • Hydrogen valence = 1 • Oxygen = 2; single and double bonds • Nitrogen = 3; single, double and triple bonds • Carbon = 4 - single, double and triple bonds

35. As you can see bystudying these two charts, there is a relationship between the number of bonding electron pairs and the number of lone electron pairs and the shape of a molecule.