2. Subatomic particles Electron Proton Neutron NameSymbolCharge Relative mass Actual mass (g) e-e- p+p+ n0n0 +1 0 1/1840 1 1 9.11 x 10 -28 1.67 x 10 -24

3. Counting the Pieces  Atomic Number = number of protons in the nucleus  # of protons determines kind of atom  The same as the number of electrons in the neutral atom.  Mass Number = the number of protons + neutrons.  These account for most of mass

4. Counting the Pieces  Protons: equal to atomic number  Neutrons: Mass Number – Atomic Number  Electrons: In a neutral atom equal to atomic number

5. Symbols  Contain the symbol of the element, the mass number and the atomic number.

6. Symbols  Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number

7. Symbols  Find the  number of protons  number of neutrons  number of electrons  Atomic number  Mass Number F 19 9

8. Symbols n Find the –number of protons –number of neutrons –number of electrons –Atomic number –Mass Number Br 80 35

9. Symbols n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

10. Symbols n if an element has 91 protons and 140 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol

11. What if Atoms Aren’t Neutral  Ions: charged atoms resulting from the loss or gain of electrons

12. What if Atoms Aren’t Neutral  Anion: negatively charged ion; result from gaining electrons  Take the number of electrons in a neutral atom and add the absolute value of the charge 81 35 Br 1- Identify: Number of Protons Number of Neutrons Number of Electrons

13. What if Atoms Aren’t Neutral  Cation: positively charged ion; result from the loss of electrons  Take the number of electrons in a neutral atom and subtract the value of the charge 27 13 Al 3+ Identify: Number of Protons Number of Neutrons Number of Electrons

14. Isotopes  Atoms of the same element can have different numbers of neutrons  Different mass numbers  Called isotopes

15. Naming Isotopes  We can also put the mass number after the name of the element.  carbon- 12  carbon -14  uranium-235

16. Atomic Mass  How heavy is an atom of oxygen?  There are different kinds of oxygen atoms  We are more concerned with average atomic mass  Average atomic mass is based on abundance of each element in nature.  We don’t use grams because the numbers would be too small

17. Measuring Atomic Mass  Unit is the Atomic Mass Unit (amu)  It is one twelfth the mass of a carbon- 12 atom  Each isotope has its own atomic mass, thus we determine the average from percent abundance

18. Atomic Mass  Is not a whole number because it is an average.  are the decimal numbers on the periodic table.

19. Modern Periodic Table  The modern periodic table consists of Rows and Columns  Rows -  Horizontal  Also known as Periods  Numbered 1-7  Columns -  Vertical  Also known as Groups and Families  Numbered 1-18

20. Metals  The most common class of elements is Metals  Metals become cations  What is a cation? How are they formed?  Positively charged atom - Lose electrons  Metals are generally solid (except Hg), conductive of heat and electricity, malleable, ductile, and shiny

21. Alkali Metals  Group 1 elements are known as Alkali Metals  Alkali metals include Li, Na, K, Rb, Cs, Fr  Alkali metals are generally dull, soft, and reactive – rarely found as free elements

22. Alkaline Earth Metals  Group 2 elements are known as Alkaline Earth Metals  Alkaline earth metals include Be, Mg, Ca, Sr, Ba, and Ra  Alkaline earth metals are harder, denser, and stronger than alkali metals  Less reactive than alkali metals, but still rarely found as free elements

23. Transition Metals  Elements in groups 3-12 (3B-2B) are known as Transition Metals  Transition metals include Mn, Fe, Ag, Au, Mo, etc.  Lanthanide and Actinide Series elements fill in the f orbitals – known as inner transition elements

24. Metalloids  Elements that border the staircase on the periodic table are known as Metalloids  Metalloids include: B, Si, Ge, As, Sb, Te, Po, At  Metalloids have properties of both metals and nonmetals

25. Nonmetals  Nonmetals are found to the right of the staircase on the periodic table  Nonmetals generally become anions  What is an Anion? How are they formed?  Negatively charged atom - Gain electrons  Nonmetals are often gases or dull, brittle solids  Nonmetals generally show poor conductivity, ductility, and malleability

26. Halogens  Group 17 elements are known as Halogens  Halogens include F, Cl, Br, and I  Halogens are the most reactive nonmetals – often found in compounds

27. Noble Gases  Elements in group 18 are known as Noble Gases  Noble Gases include He, Ne, Ar, Kr, Xe, Rn  Noble gases are extremely unreactive

28. Legend

30. Pure Substances  Cannot be physically separated  Every sample has the same characteristics and they can be used to identify a substance

31. Elements  Are made up of ONE type of atom  Atoms are the smallest unit of an element that maintains the chemical identity of that element  They can be found on the Periodic Table  Examples: Carbon, Nitrogen, Calcium

32. Compounds  Can be broken down into simple stable substances  Are made up of two or more types of atoms that are chemically bonded  Examples: Water (H 2 O), sugar (C 12 H 22 O 11 )

33. Mixtures  A blend of two or more kinds of matter, each which retains its own identity and properties

34. Homogeneous Mixtures  Have uniform composition  Also known as SOLUTIONS  Examples: salt water, tea

35. Solutions  ALLOYS are solid solutions that contain at least 1 metal  They are blended together so that they have more desirable properties  Some alloys you may know are:  Stainless Steel: iron, chromium, and zinc  Brass: zinc and copper  Bronze: tin and copper  Sterling Silver: copper and silver

36. Heterogeneous Mixtures  Do not have uniform composition  You can see the particles in them  Examples:  Sand on the beach (contains sand, shells, rocks, bugs, etc)  Soil (contains dirt, rocks, worms, etc)  Chicken Soup (contains water, chicken, veggies etc)

37. Suspensions  A heterogeneous mixture where the solid particles eventually settle out of solution  Examples:  Muddy water  Mixtures of two solids  Paint

38. Properties of Matter  All pure substances have characteristic properties  Properties are used to distinguish between substances  Properties are also used to separate substances

39. Physical Properties  A Physical Property is a characteristic that can be observed or measured without changing the composition of the substance  Physical properties describe the substance itself  Examples  Physical State  Color  Mass, shape, length  Magnetic properties

40. Chemical Properties  A Chemical Property indicates how a substance will react with another  Chemical properties cannot be determined without changing the identity of the substance  Examples:  Iron Rusting  Silver Tarnishing

41. Physical Changes  A Physical Change is a change in a substance that does not alter the substance’s identity  Examples:  Grinding  Cutting  Melting  Boiling

42. Chemical Changes  A change in which one or more substances are converted into different substances is called a Chemical Change  Signs of a Chemical Change:  Color Change  Gas is Released  Temperature Change  Precipitate – Solid falls out of solution  Substance Disappears

43. Electrons  Electrons fill in an atom in energy levels  Electrons occupy the LOWEST available energy level  Energy Levels hold limited amounts of electrons  1 st Level – 2 electrons  2 nd Level – 8 electrons  3 rd Level – 18 electrons  4 th Level – 32 electrons

44. Chlorine (Cl) P = 17 N = 18 E = 17

45. Nitrogen (N) P = 7 N = 7 E = 7

46. Aluminum (Al) P = 13 N = 14 E = 13

47. Valence Electrons  Electrons in outermost shell that determine chemical behavior  Maximum of 8 valence electrons  Atoms with same valence electrons will act similarly  Group 1 elements?  1 valence electron  Group 17 elements?  7 valence electrons

48. How Atoms Combine  Two or more atoms that are chemically combined make up a compound  The combination results in a chemical bond, a force which holds elements together in a compound

49. Covalent Bonds  Covalent Bonds are formed when atoms in a compound share electrons  Molecule – two or more atoms held together by a covalent bond  Usually occurs between nonmetals

50. Covalent Bonding in Water

51. Ions  An atom that has gained or lost an electron is called an ion.  Multiple atoms can combine to form an ion – called a Polyatomic Ion  Silicate (SiO 4 4- ) and Carbonate (CO 3 2- ) are important in forming materials at Earth’s Surface

52. Ionic Bonding  Positive and negative ions attract each other  Ionic Bonds occur when oppositely charged ions form a compound  Usually consist of 1 metal and 1 nonmetal  Positive ion written first in chemical formula (NaCl)  Ionic compounds have a neutral charge

53. Ionic Bonding in NaCl

54. Metallic Bonds  Metals share valence electrons between all atoms  Like a group of positive ions in a sea of electrons

55. Acids and Bases  An Acid is a substance that produces Hydrogen Ions in water (H + )  Acids:  Sting to the touch  Taste Sour  React with metals

56. Acids and Bases  A Base is a substance that produces hydroxide ions (OH - ) in water  Bases  Are slippery to the touch  Taste Bitter  Do not react with metals

57. pH Scale  Measures the amount of hydrogen ions in a solution  0 - 6 Acidic  7 = Neutral  8 - 14 Basic

58. Indicators Indicators are substances that turn colors at different pH levels Examples: Litmus Phenolphthalein Base indicator Universal Indicator: ACID NEUTRAL BASE Cabbage Juice ACID NEUTRAL BASE