2. 2.1 Periodic Table Atomic Structure Trends Bonding

3. Atoms An atom is the smallest part of an element that retains all the properties of that element. Atoms are made up of three types of particles –a nucleus consisting of protons and neutrons –electrons in shells around the nucleus. The most common way to picture an atom is the solar system model (Bohr model). The number of protons in the nucleus determines the element, as well determining the number of electrons.

4. Atomic structure: Nucleus Contains two kinds of particles that are approximately the same size and mass –positively charged protons –uncharged neutrons. Since like charges repel, there is some sort of "nuclear glue" that holds the nucleus together. The atomic number of an element is the number of protons in the nucleus of an atom of that element. + Proton Neutron

5. Atomic Structure: Electrons Electrons are negatively charged Electrons are smaller and have less mass than protons and neutrons Electrons orbit the nucleus Atoms are usually neutral so they have the same number of protons and electrons (charges cancel each other out) When there are a different number of electrons in orbit than there are protons in the nucleus, the atom is called an ion because it carries an electrical charge

6. Atomic Structure: Electron shells Electrons are arranged in energy levels (shells) or orbits around the nucleus. There is a definite arrangement of the electrons in these shells and a maximum number of electrons possible in each shell. –The maximum number of electrons in the first energy level is 2. –After the first energy level is filled, the second starts filling up, according to the number of positive charges in the nucleus. The maximum number of electron in the second energy level is 8 electrons. –Then the third shell starts to fill, and so on http://www.stillwaterpalladium.com/periodic.html Palladium

7. Electron Energy Levels cHolt, Rinehart & Winston, Science Spectrum

8. Atoms summary The nucleus contains two particles –Positively charged protons –Neutral neutrons –Roughly the same size Negatively charged electrons orbit the nucleus –Smaller in size and mass than protons & neutrons –Arranged in energy shells that have rules about the way they fill up + Proton Neutron - Electron

9. Ions Ions are atoms that have an unequal number of electrons and protons, so they have a charge associated with them The amount of charge depends upon how much difference in numbers of protons and electrons associated with the ion Positive ions are cations (Na + ) Negative ions are anions (I - ) http://www.cybered.net

10. Element Composed of identical atoms Cannot be broken down into simpler substances. Properties are due to atomic structure of element Atoms make up matter and atoms make up elements. There are 111 known elements, each has its own abbreviation Examples: –Copper –Carbon http://www.webelements.com/webelements/elements/text/C/key.html http://www.webelements.com/webelements/elements/text/Cu/key.html

11. http://cougar.slvhs.slv.k12.ca.us/~pboomer/chemlectures/periodictable.jpg All About the Periodic Table

12. Elements & the Periodic Table Elements are organized in the table by atomic structure and properties They are listed in order by atomic number In order to use & understand the table, you need to know how to read the information it contains. www.cybered.net

13. Reading the Periodic Table http://www.classzone.com/books/earth_science/terc/content/investigations/es0501/es0501page06.cfm

14. Complete the following chart indicating the number of protons, neutrons and electrons each atom has.

15. ElementProtonsNeutronsElectrons Chlorine-35 Iron-56 Magnesium-24 Lead-208 Nitrogen-14 17 18 263026 12 82126 82 77 7

16. Parts of the Periodic Table Element’s info Period- horizontal row Group (sometimes called a family) -vertical row -similar physical and chemical properties Group number is at top of group Some tables show groups labeled as 1 – 18 from left to right

17. Arrangement of the Periodic Table The elements in the following Periodic Table are arranged in rows (Periods) according to atomic number And in columns (Groups) according to the configuration of the outer orbit (energy level). If you go along a period from left to right, the elements are numbered 1 - H, 2 - He, 3 - Li, 4 - Be, 5 - B, and so on. This atomic number is also the number of protons in the element's nucleus. If you go down a group, each element has the same number of electrons in its outer orbit or shell. –H, Li, Na, K and so on, all have one electron in the outer energy level. –O, S, Se, etc. all have 6 electrons in the outer orbit. The number of electrons in the outer energy level determines the element's chemical properties.

18. are periodic functions of the number of valence electrons an atom has. Groups or families (vertical column) of elements are listed by increasing atomic number and they have similar chemical properties. Periods (horizontal rows) of elements usually start with a reactive solid and end with an unreactive gas, called a noble gas. Periods also increase by atomic number.

19. Element Properties and the Periodic Table Metals, nonmetals and Metalloids (semimetals) –Elements on the left side of the periodic table are metals. solid substances that are good conductors of heat and electricity. can be easily formed into many shapes. –Elements B, Si, Ge, As, Sb, Te, and At are metalloids. conduct heat and electricity better than nonmetals, but not as well as metals easier to shape than nonmetals, but not as easy as metals. All are solid at room temperature. –Elements on the right of the periodic table are Nonmetals. poor conductors of heat and electricity not easily formed into shapes. Most are gases at room temperature, except for C, P, S, Se and I, which are solids and Br, which is a liquid. H is also a nonmetal.

20. Metals metals  metalloids  nonmetals Properties of Metals: Luster Ductile Malleable Good conductors of heat and electricity, Lose electrons to form + ions Metals

21. Non-metals metals  metalloids  nonmetals Nonmetals Properties: dull brittle Insulators (poor conductors) gain electrons to form – ions Non-metals

22. Metalloids metals  metalloids  nonmetals Metalloids Metalloid (Semimetals) Border the diagonal line between metals and nonmetals. Have some properties of both metals and nonmetals. Act as a metal with a nonmetal or as a nonmetal with a metal.

23. Noble Gases are non- reactive, they have full electron shells

24. It’s the valence electrons orbiting in the outer energy level that allow one atom to interact with other atoms so they can be linked together.

26. Valence Electrons 1 valence electron 2 valence electrons 3 valence electrons 4 valence electrons 5 valence electrons 6 valence electrons 7 valence electrons 8 valence electrons Electrons in the outermost level are called valence electrons.

27. Periodic Table: Predicting Reactions Reactions occur to complete orbital shells Elements will lose or gain electrons to have eight electrons (same as noble gas) in their outer energy level. As you know, atoms "like" to have their outer orbit completely filled or completely empty. (Only for TAKS) –Hydrogen has 1 more electron than having its orbit empty. –Since Boron (B) already has its first orbit full and will allow a maximum of 8 electrons in its second orbit, +3 indicates it has 3 more electrons than having its outer orbit empty. –Likewise, Chlorine (Cl) is one electron short (-1) of filling up its outer orbit or energy level. Groups are or vertical columns and are now numbered 1-18. All elements in the same group have the same number of valence electrons (electrons in the outer energy level) and react similarly.

28. Chemical Bonding Chemical Bonding - combining atoms of elements to form new substances (compounds) Valance electrons - the electrons in the outermost energy level. The maximum number is 8. These are the electrons involved in chemical bonding Atoms give up or take on electrons in order to have a full shell. This is why chemical bonding occurs.

29. Ionic Bonding Lose electrons (+) Accept electrons (-) Occurs typically between a metal and a non-metal. Metal atom loses electrons which are taken by the non-metal. Metal ion produced has a positive charge; the non-metal ion is negatively charged.

30. Ionic Bonding Ionic Bonding – Between a metal and a nonmetal atom Bonding that involves a transfer of electrons Formation of ions or "charged“ atoms One atom gains electrons and the other atom loses electrons Positive ions attract negative ions and form ionic bonds.

31. Covalent Bonding Covalent Bonds -Between Nonmetal Atoms Covalent bonds are forces of attraction formed when atoms of a molecule share electrons. The term sharing electrons indicates that the valance electrons of the atoms become part of the orbitals of more than one atom of the molecule.

32. Types of Covalent Bonds

33. Metallic Bonds Metallic Bonds –Only Metal Atoms bonds formed by the atoms of metals, in which the outer electrons of the atoms form a common electron cloud or "sea" of electrons. The electron cloud allows charge to flow, which is why metals are conductors The positive nuclei of metal atoms are surrounded by free-moving or mobile electrons that are attracted by the nuclei Only Metal Atoms bond this way! Positive metallic ions Electron cloud formed by extra electrons

34. Predicting Ionic Charges Group (IA) 1: Lose 1 electron to form 1+ ions H+H+H+H+ Li + Na + K+K+K+K+

35. Predicting Ionic Charges Group (IIA) 2: Loses 2 electrons to form 2+ ions Be 2+ Mg 2 + Ca 2+ Sr 2 + Ba 2+

36. Predicting Ionic Charges Group (IIIA) 13: Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

37. Predicting Ionic Charges Group (IVA) 14: Lose 4 electrons or gain 4 electrons? Many Group IVA (14) elements rarely form ions. Many Group IVA (14) elements rarely form ions.

38. Predicting Ionic Charges Group (VA) 15: Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

39. Predicting Ionic Charges Group (VIA) 16: Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

40. Predicting Ionic Charges Group (VIIA) 17: Gains 1 electron to form 1- ions F 1- Cl 1- Br 1- Fluoride Chloride Bromide I 1- Iodide

41. Predicting Ionic Charges Group (VIIIA) 18: Stable Noble gases do not form ions!

42. Predicting Ionic Charges Groups 3 - 12: Many transition elements Many transition elements have more than one possible oxidation state. Roman numerals are used to indicate the oxidation state Iron(II) = Fe 2+ Iron(III) = Fe 3+

43. Predicting Ionic Charges Groups 3 - 12: Some transition elements Some transition elements have only one possible oxidation state. have only one possible oxidation state. Zinc = Zn 2+ Silver = Ag +

44. Writing Ionic Compound Formulas Example: Iron(III) with chloride 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 3 Fe 3+ Cl - Balanced! Iron(III) chloride 4. Name the ions

45. Barium with nitrate Example: Barium with nitrate 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2 Ba 2+ NO 3 - Balanced! Barium nitrate 4. Name the ions

46. Example: Ammonium with sulfate 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2 NH 4 + SO 4 2- Your Turn! Ammonium sulfate 4. Name the ions

47. Example: Aluminum with sulfide 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! 2 3 Al 3+ S 2- Aluminum sulfide 4. Name the ions

48. Magnesium with carbonate Example: Magnesium with carbonate 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced! Magnesium carbonate 4. Name the ions

49. Example: Zinc with hydroxide 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Not balanced! ( ) 2 Zn 2+ OH - Zinc hydroxide 4. Name the ions

50. Example: Calcium with phosphate 1. Write the formulas for the cation and anion, including CHARGES! Ca 2+ PO 4 3- 2. Check to see if charges are balanced. Not balanced! Calcium phosphate Ca 2+ PO 4 3- 23 ( ) 4. Name the ions 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion.

51. Chemical Formulas and Chemical Compounds Heart cell rhythm depends on the opening and closing of a complex series of valves on the cell membrane, called ion channels. Some valves let certain ions like potassium (K+) flow out, others let different ions like sodium (Na+) flow in. There are also pumps that actively move ions one direction or another.