Atoms and Bonding. Chapter 18 Physical Science

Transcription

1 Atoms and Bonding Chapter 18 Physical Science

2 Atoms and Bonding: Chemical Bonding The combining of atoms of elements to form new substances. Bonding of atoms determine a compound s properties. Compounds More than one kind of atom chemically combined. Atoms combine in distinct ways based on their structure (valence electrons).

3 Lewis-dot Diagram Bonds can be shown using a Lewis-dot diagram. The diagram uses dots that represent the outer electrons of each atom involved in the bonding. ulyopnxjaz8

4 Electrons & Energy Levels: Atoms are by nature neutral. i.e. Electrons will equal protons. Electrons are held by electromagnetism to the protons in the nucleus. Electrons are arranged around the nucleus in an electron cloud. Electron clouds are made up of energy levels

5 Electrons & Energy Levels: Electron clouds are made up of energy levels. Maximum electrons per energy level (2n 2 ) 1 st = (2)(1 2 ) = 2 nd = (2)(2 2 ) = 3 rd = (2)(3 2 ) = 4 th = (2)(4 2 ) = 5 th = (2)(5 2 ) = or 18 8 or 32 8 or 50 *Note: No element known today has attained an outer energy level of 50 electrons.

6 Electrons & Energy Levels: When an atom s outer energy level is completely filled with electrons, it becomes stable (unreactive). Q: Which elements in the periodic table have complete outer energy levels? A: Noble Gases (Group 8a) Q: Based on this principle, would you rather fly in a hydrogen filled or helium filled airship?

7 The Octet Rule: Most elements with more than 2 energy levels are stable with 8 outer electrons even though they can hold 18, 32, or 50 electrons. This rule applies to the behavior of most elements and compounds, but is not universal.

8 The Octet Rule, cont d: If an atom has only one outer electron (valence electron), it will likely lose that electron to another atom. With this electron missing, the element becomes stable with its next inner energy level full, but is now an ion. Specifically a + ion (cation).

9 Ionic Bonding: Ionic bonds involve the transfer of electrons from one atom to another. Always formed between a metal and a non-metal. An atom in an ionic bond will become either positive or negatively charged.

10 Ionic Bonding: Ion (charged particle) = An atom that has lost or gained electrons. Becoming ions for an ionic bond: Sodium (Na) Fluorine (F) 11P 9P 11 Protons (+11) 11 Electrons + (-11) net charge = 0 9 Protons (+9) 9 Electrons + (-9) net charge = 0

11 Ionic Bonding: Since Sodium becomes positively charged and Fluorine becomes negatively charged, electromagnetism holds the atoms together to form a compound Sodium fluoride (NaF) Sodium (Na +1 ) Fluorine (F -1 ) 11P 9P 11 Protons (+11) 11 Electrons + (-10) net charge = +1 9 Protons (+9) 9 Electrons + (-10) net charge = -1

12 Electron Affinity: Electron affinity is the tendency of an atom to attract electrons. Q: Which element has the highest electron affinity, calcium or chlorine? A: Chlorine

13 Ionization Energy: Ionization energy is the energy required for ionization. Ionization involves overcoming the attraction of electrons to protons. Ionization energy is proportional to the number of valence electrons. i.e. Atoms with tightly held valence electrons have high ionization energy & high electron affinity.

14 Ionic Bonds cont d: Since (+) ions are attracted to (-) ions and like ions repel, ionic compounds tend to form crystal lattices a repeating arrangement of ions. This arrangement gives rise to stability. i.e. A high amount of energy is required to overcome ionic attractions.

15 Ionic Bonds cont d: Ionic bonds can be shown using Lewis-Dot Diagrams Dots represent outer electrons (valence electrons) Sodium Chloride

16 Covalent Bonds: Covalent bonds involve the sharing of valence electrons. Most likely to occur between atoms w/high ionization energy & electron affinity. - i.e. Atoms in covalent bonds do not lose electrons easily & tend to attract electrons - The positive nuclei of two atoms simultaneously attract each others electrons.

17 Covalent Bonds: Covalent bonds can be shown using Lewis- Dot Diagrams Dots represent outer electrons (valence electrons) e.g. Diatomic Hydrogen Diatomic Iodine

18 Covalent Bonds: Hydrogen Chloride Water Ammonia

19 Covalent Bonds: Covalently bonded atoms are called molecules. - Molecule smallest particle of a covalently bonded substance that has all the properties of that substance. - Molecules tend to have low melting points or are gases. - i.e. Molecules have weak attractions to one another Exceptions are network solids like diamond, graphite, quartz - In network solids there is no separation of molecules it is one big molecule (macromolecule).

20 Covalent Bonds: Polyatomic Ions Covalently bonded atoms that have a charge and act like a single atom. - These tend to form ionic bonds w/other atoms - Ammonium NH Phosphate PO Sulfite SO 3-2

21 Covalent Bonds: Polar Covalent Bonds Unequal sharing of electrons e.g. Water (H 2 O) - O H + + H

22 Covalent Bonds: H O H Oxygen tends to possess the Hydrogens electrons more often; Hydrogen becomes positively charged & Oxygen becomes negatively charged. This is why water molecules adhere to each other.

23 Metallic Bonds: Metallic bonds are bonds between two metals. Metals tend to give up electrons. Positive nuclei of atoms of metals are surrounded by free-moving electrons. i.e. They form a common electron cloud & all equally share their outer electrons.

24 Metallic Bonds: Free floating electrons allow for metallic properties because: - Atoms can slide by each other when twisted or hammered. - Electrons can flow easily (high conductivity). - The overall attraction of electrons give metals a high melting point.

25 Predicting Bond Types: Metal & Non-Metal = Ionic Bonded Compounds Non-metal & Non-metal = Covalently Bonded Compounds Metal & Metal = Metallic Bonds

26 Formation of Water: H 2H 2 H H H Hydrogen + Oxygen Reactants O O 2H 2 O Water O H H O 2 O H H Products

27 Oxidation Numbers: Oxidation numbers are the number of electrons an atom gains, loses, or shares in a chemical bond. e.g. Sodium (Na) = 1 valence electron = Na+ when bonding occurs = +1 oxidation number Magnesium (Mg) = +2

28 Oxidation Numbers: Q: What is the oxidation number of Fluorine? A: -1 Q: Carbon? A: +4 or -4

29 Using Oxidation Number to predict Chemical Bonding and Formulas. - Step 1: Identify the element symbol or the polyatomic ion formula for each part of the compound name. - Step 2: Determine the oxidation number for each part. - Step 3: The sum of oxidation numbers must equal zero for the completed formula. - Step 4: If the sum is not zero, use Least Common Multiple of the oxidation numbers to determine how many of each element or ion is needed so that the sum will be zero.

30 Ex: Write the formula for Sodium chloride and for Calcium chloride Sodium chloride step 1: Na Cl step 2: +1-1 step 3: (+1) + (-1) = 0 Formula = NaCl Calcium chloride step 1: Ca Cl step 2: +2-1 step 3: (+2) + (-1) 0 step 4: 1(+2) + 2(-1) = 0 Step 4 tells me I need 1 Ca and 2 Cl: Formula = CaCl 2

31 Oxidations of Common Elements: +1 Ammonium NH 4 + Cuprous Cu + Hydrogen H + Mercurous Hg + Potassium K + Silver Ag + Sodium Na Barium Ba ++ Aluminum Al +++ Calcium Ca ++ Arsenic As +++ Cupric Cu ++ Chromic Cr +++ Ferrous Fe ++ Ferric Fe +++ Lead Pb ++ Magnesium Mg ++ Manganese Mn ++ Mercuric Hg ++ Nickel Ni ++ Tin Sn ++ Zinc Zn ++

32 Oxidations of Common Elements: -1 Acetate C2H3O2 - Bicarbonate HCO3 - Bisulfate HSO4 - Bromate BrO3 - Bromide Br - Chlorate ClO3 - Chloride Cl - Cyanide CN - Flouride F - Hydroxide OH - Iodide I - Nitrate NO3 - Nitrite NO2 - Permanganate MnO Carbonate CO 3 -- Chromate CrO 4 -- Dichromate Cr 2 O 7 -- Oxide O -- Peroxide O 2 -- Sulfate SO 4 -- Sufide S -- Sulfite SO Nitride N --- Phosphate PO 4 ---