1. What is an atom? (c) McGraw Hill Ryerson 2007

2. 4.1 Atomic Theory and Bonding An atom is the smallest particle of an element that still has the properties of that element  50 million atoms, lined up end to end = 1 cm  An atom = proton(s) + neutron(s) + electron(s) Atoms join together to form compounds.  A compound is a pure substance that is composed of two or more atoms combined in a specific way.  Oxygen and hydrogen are atoms/elements; H 2 O is a compound. A chemical change occurs when the arrangement of atoms in compounds changes to form new compounds. See pages 168 - 169

3. (c) McGraw Hill Ryerson 2007 Atomic Theory Atoms are made up of smaller particles called subatomic particles. The nucleus is at the center of an atom.  The nucleus is composed of protons and neutrons.  Electrons exist in the space surrounding the nucleus called shells. See page 170

4. (c) McGraw Hill Ryerson 2007

5.  Protons-positive (1+) electric charge  Neutrons-no charge  Electrons-Negative (1-) electric charge  # of protons = # of electrons in every atom  Nuclear charge = charge on the nucleus = # of protons  Atomic number = # of protons = # of electrons (c) McGraw Hill Ryerson 2007

6. Periodic Table and Ion Formation Atoms gain and lose electrons to form bonds.  The atoms become electrically charged particles called ions.  Metals lose electrons and become positive ions (cations).  Some metals (multivalent) lose electrons in different ways.  For example, iron, Fe, loses either two (Fe 2+ ) or three (Fe 3+ ) electrons  Non-metals gain electrons and become negative ions (anions).  Atoms gain and lose electrons in an attempt to have the same number of valence electrons (electrons farthest from the nucleus) as the nearest noble gas in the periodic table. See page 173 ~ ~

7. Ion Summary MetalsNon-metals Lose electronsGain electrons Become positively charged Become negatively charge

8. Let’s get familiar with each box on the Periodic Table Atomic Number Chemical Symbol Chemical Name Atomic Mass 11 Na Sodium 23.0

9. Let’s look at the Atomic # On our Periodic Table, it’s the # on the TOP 11 Na Sodium 23.0 Think APE!!! Atomic number = Protons = Electrons

10. At this point we know A=P=E Atomic Number = Protons= Electrons Let’s Practice!!!

11. LET’S PRACTICE…THINK A=P=E 1.Atomic number? 2.Number of Protons? 3.Number of Electrons? 8 O OXYGEN 16 8 8 8

12. 11 Na Sodium 23.0 On our Periodic Tables, it’s the number on the bottom. Notice the # is in decimal form. For our purposes, we will round to the nearest whole # Atomic Mass

13. This is what we know: 1. Atomic # (which also tells us the # of Protons & Electrons) 2. Atomic Mass SO… how do we find the # of Neutrons?

14. M-A=N Mass- Atomic Number = Neutrons

15. Let’s figure it out! What ‘s the Atomic #? How many Protons are there? How many Electrons are there? What is the Atomic mass? What part of the Atomic Mass represents the protons? So, if we have 11 protons in the Atomic Mass, how much mass then must represent the Neutrons? 11 Na Sodium 23.0 11 23 11

16. How do we find Neutrons?? M-A= N 23 - 11 12 Neutrons

17. Think APE MAN Find the number of protons, neutrons, and electrons in Fluorine!! A=P=E 9 protons, 9 electrons M-A=N 19-9 = 10 Neutrons Practice Time: with a partner go through the chart on the next page (fill out).

18. Chem. Name Chem. Symbol Atomic Number Atomic Mass # of Protons# of Neutrons # of Electrons K 50 36 195 Barium

19. (c) McGraw Hill Ryerson 2007 Bohr Diagrams Bohr diagrams show how many electrons appear in each electron shell around an atom.  Electrons in the outermost shell are called valence electrons.  Think of the shells as being 3-D like spheres, not 2-D like circles. See page 174 What element is this? It has 2 + 8 + 8 = 18 electrons, and therefore, 18 protons. It has three electron shells, so it is in period 3. It has eight electrons in the outer (valence) shell. 18 p 22 n argon

20. Practice  Draw bohr diagrams for Mg, Ca, & O  (draw long hand & short hand on the board) Worksheet on Bohr models (c) McGraw Hill Ryerson 2007

21. Patterns of Electron Arrangement in Periods and Groups Electrons appear in shells in a very predictable manner. There is a maximum of two electrons in the first shell, eight in the 2nd shell, and eight in the 3rd shell.  The period number = the number of shells in the atom.  Except for the transition elements, the last digit of the group number = the number of electrons in the valence shell. See page 175  The noble gas elements have full electron shells and are very stable.

22. Practice problems  Pg 176  Identify # of occupied shells for  Calcium, Ca  Krypton, Kr  Sulfur, S  Iodine, I  Identify # of valence electrons  Chlorine, Cl  Magnesium, Mg  Strontium, Sr  Bromine, Br (c) McGraw Hill Ryerson 2007

23. Forming Compounds When two atoms get close together, their valence electrons interact.  If the valence electrons can combine to form a low-energy bond, a compound is formed.  Each atom in the compound attempts to have the stable number of valence electrons as the nearest noble gas. Forming Compounds Metals may lose electrons to other atoms forming cations Non-metals may gain electrons from other atoms forming anions Also atoms may share electrons (covalent bond). See pages 176 - 177

24. Ionic Bonds Ionic bonds form when electrons are transferred from positive ions to negative ions.  Must contain a positive (usually a metal) and a negative ion (usually a non metal) Ionic bonding is when one or more electrons transfer from one metal to a non metal Lets go through some examples on the board. Start by drawing out the bohr models  CaF 2  NaCl (c) McGraw Hill Ryerson 2007

25. Forming Compounds (continued) Ionic bonds are formed between positive ions and negative ions.  Generally, this is a metal (+) and a non-metal (-) ion.  For example, lithium and oxygen form an ionic bond in the compound Li 2 O. See pages 176 - 177 lithium oxygen + Electrons are transferred from the positive ions to negative ions Li + O 2- Li + lithium oxide, Li 2 O

26. Covalent bonds form when electrons are shared between two non-metals.  Electrons stay with their atom but overlap with other shells. Bonding pair- pair of electrons involved in a covalent bond Lone pair- a pair of electrons in the valence shell that is not used in bonding Lets do some bohr diagram examples on the board  HF  H 2 O  NH 3 (c) McGraw Hill Ryerson 2007

27. Lewis Diagrams Lewis diagrams illustrate chemical bonding by showing only an atom’s valence electrons and the chemical symbol.  Dots representing electrons are placed around the element symbols at the points of the compass (north, east, south, and west).  Electron dots are placed singly until the fifth electron is reached then they are paired. See page 178

28. (c) McGraw Hill Ryerson 2007 Lewis Diagrams of Ions Lewis diagrams can be used to represent ions and ionic bonds.  For positive ions, one electron dot is removed from the valence shell for each positive charge.  For negative ions, one electron dot is added to each valence shell for each negative charge.  Square brackets are placed around each ion to indicate transfer of electrons. See page 179 BeCl Each beryllium has two electrons to transfer away, and each chlorine can receive one more electron. BeCl BeCl Since Be 2+ can donate two electrons and each Cl – can accept only one, two Cl – ions are necessary. beryllium chloride 2+ – –

29. (c) McGraw Hill Ryerson 2007 Lewis Diagrams of Covalent Bonds Lewis diagrams can also represent covalent bonds.  Like Bohr diagrams, valence electrons are drawn to show sharing of electrons.  The shared pairs of electrons are usually drawn as a straight line. See page 179

30. (c) McGraw Hill Ryerson 2007 Lewis Diagrams of Diatomic Molecules Diatomic molecule- molecules composed of only two atoms which are the same. - H 2 or O 2 See page 180 O O O O O O Several non-metals join to form diatomic molecules. Valence electrons are shared, here in two pairs. This is drawn as a double bond. Diatomic molecules, like O 2, are also easy to draw as Lewis diagrams. Take the Section 4.1 Quiz