3. Section 6.1 Atoms and Moles B. Atomic Masses: Counting Atoms by Weighing Atoms have very tiny masses so scientists made a unit to avoid using very small numbers. 1 atomic mass unit (amu) = 1.66  10 -24 g The average atomic mass for an element is the weighted average of the masses of all the isotopes of an element.

4. Section 6.1 Atoms and Moles C. The Mole One mole of anything contains 6.022 x 10 23 units of that substance. –Avogadro’s number is 6.022 x 10 23. The mole is defined as the number of atoms in exactly 12g of carbon-12.

5. Section 6.1 Atoms and Moles A. Molar Mass A compound is a collection of atoms bonded together. The molar mass of a compound is obtained by summing the masses of the component atoms.

6. Section 6.1 Atoms and Moles B. Percent Composition of Compounds Percent composition consists of the mass percent of each element in a compound: Mass percent =

7. Section 6.1 Atoms and Moles B. Calculation of Empirical Formulas

8. Section 6.1 Atoms and Moles Find the Empirical formula 25.95% Nitrogen 74.06% Oxygen

9. Section 6.1 Atoms and Moles B. Mole-mole Relationships A balanced equation can predict the moles of product that a given number of moles of reactants will yield. The mole ratio allows us to convert from moles of one substance in a balanced equation to moles of a second substance in the equation.

10. Section 6.1 Atoms and Moles 9.2 B. Mass Calculations Using Scientific Notation Stoichiometry is the process of using a balanced chemical equation to determine the relative masses of reactants and products involved in a reaction. –Scientific notation can be used for the masses of any substance in a chemical equation. To solve: gA -> mol A -> mol B -> g B To convert between moles and grams we use the molar masses of the substance. To convert between moles we use the mole ration from the balanced equation.

11. Temperature7a Temperature is a measure of the average kinetic energy of the molecules (/atoms) of a substance. ◦ In a hot sample, the molecules are moving much faster than in a cold sample. What is temperature?

12. Heat Heat is energy transferred from molecules at a higher temperature to molecules at a lower temperature. What happens when you heat up a substance?

13. Heat Flow Problems To calculate heat, use the following formula: ◦ Energy = mass * specific heat * temperature change Always given to you!

14. Latent Heat7d cont Latent (hidden) heat is energy added to a substance that doesn’t change the temperature. ◦ Instead, it is used to change the phase solid  liquidliquid  gas What is latent heat and what is it used for?

15. Heat added here doesn’t change the temperature of the substance, it is causing the phase change.

16. ΔH – Heat of Reaction7b cont Endothermic ◦ The products have more energy than the reactants. ◦ Heat must be put into the reaction, so ΔH is positive.

17. Exothermic ◦ The products have less energy than the reactants. ◦ Heat must be released from the reaction, so ΔH is negative.

18. Summary What is temperature What happens when you heat a substance? Describe endo/exothermic reactions. What is ΔH for endo/exothermic reactions? Where does the energy go during phase changes? Be able to calculate heat problems. What is latent heat used for?

20. All digits in a number are significant except: 1. 0s to the left of ALL the other digits. 2. 0s to the right of ALL the other digits when there is no decimal place. -Why aren’t these 0s significant? Because they weren’t measured, they are just there to tell you what place the first number is in. a.1457 4 significant figures b.0.0025 2 significant figures c.1.008 4 significant figures d. 100 1 significant figure e. 100. 3 significant figures f. 120.0 4 significant figures Examples

21. Section 5.1 Scientific Notation and Units A. Scientific Notation Representing Large Numbers Representing Small Numbers 0.000167To obtain a number between 1 and 10 we must move the decimal point. 0.000167 = 1.67  10 -4

22. Section 5.1 Scientific Notation and Units B. Units Units provide a scale on which to represent the results of a measurement.

23. C. Density Density is the amount of matter (atoms) present in a given volume of substance.

24. The Scientific Method 1.Observation 2.Ask a Question 3.Form a Hypothesis 4.Set up an Experiment 5.Record Data 6.Draw a Conclusion 7.Repeat…..Repeat…..Repeat 8.Theory…. maybe

25. Hypothesis vs. Theory  Hypothesis – A proposed scientific explanation for a set of observations  Theory – A well-tested explanation that unifies a broad range of observations  Plate Tectonics How does a hypothesis become a theory??

26. Theory vs. Law Theories Do Not Become Laws A natural law is a summary of behavior, it tells what happens. A theory is our attempt to explain why it happens.

27. A physical property is any characteristic of a material that can be observed without changing the material, such as color, length, or shape. Substances have physical properties that can be described and physical changes that can be observed. A chemical property is the ability or inability of a substance to combine with or change into one or more new substances. In a chemical change, the properties that give a substance its identity change.

28. A physical change is any change in the size, shape, or state of matter in which the identity of the substance is not changed. Dissolving is mixing a substance into another substance to form a solution. Mixing is a physical change in which neither substance dissolves into the other. Chemical changes change one substance into another substance. Usually chemical changes cannot be easily reversed. Forming New Substances All chemical changes produce substances that are different from the starting substances. The Atoms, molecules or bonds have been rearranged!!!

29. Elements An element is a pure substance made from atoms that all have the same number of protons. Atoms of a particular element always have the same number of protons. The number of protons in an atom of an element is the element’s atomic number. Review, don’t copy

30. Compound A compound is a distinct, pure substance made of 2 or more elements in the same, fixed ratio. Compounds are not a mixture of elements! They have their own properties and can only be broken down into elements by chemical processes. (Think water!) Compounds are written with a chemical formula showing the type and number of atoms present. E.g. H 2 O, C 6 H 12 O 6.

31. Atomic Number (Z) The number of protons. The atomic number is also equal to the number of electrons in a neutral atom (not an ion). The elements on periodic table are ordered by atomic number.

32. Mass number (A) The sum of the number of protons and neutrons in the nucleus is called the mass number. Atomic mass unit (amu) is used to measure atomic mass. –Protons & neutrons each have an AMU of 1 –Electrons have so little mass that their mass is negligible and is ignored.

33. Isotopes Atoms of the same element always have the same number of protons, but they may have different numbers of neutrons.

34. Ex. Isotope Of Neon

35. Ex. Carbon Isotopes The average atomic mass of an element is the weighted average mass of the mixture of an element’s isotopes.

36. Thomson’s Experiments (cont.) Opposite charges attract each other. Thomson concluded the cathode ray must have a negative charge and named the particles electrons. 4.2 Discovering Parts of the Atom

37. Discovering the Nucleus In Rutherford’s gold foil experiment, particles were shot through a thin sheet of gold into a detector behind the foil. 4.2 Discovering Parts of the Atom

38. Copy this slide Rutherford’s Conclusions Ernest Rutherford showed that atoms have internal structure. –A nucleus that is very small with a positive chare and a large mass. –Since the atom is mostly empty space he proposed that electrons move around the nucleus.

39. Section 3.4 Using the Periodic Table

40. Section 3.4 Using the Periodic Table B. Natural States of the Elements Diatomic Molecules Nitrogen gas contains N 2 molecules. Oxygen gas contains O 2 molecules.

41. Section 3.4 Using the Periodic Table B. Natural States of the Elements Diatomic Molecules

42. Review  Valence electrons are electrons that are found in the outer shell of the atom.  Remember that the Group Number relates to the number of Valence electrons  All atoms want a full valence shell (8 electrons). This is called the octet rule.  Electronegativity-ability to attract e-(how much an atom wants an e-)

43. Patterns (cont.) Ion Charges and the Periodic Table

44. The Sodium and Chlorine Ion are held together by an electrostatic force called an Ionic Bond.

45.  A covalent bond is a shared pair of electrons between two atoms. (nonmetals)  Atoms with similar electronegativities share the electrons fairly equally.  2+ covalently- bonded atoms make a molecule

46.  When atoms of different electronegativities bond, they don’t share electrons equally.  This unequal sharing of electrons results in a polar covalent bond. How are polar covalent bonds different from regular (nonpolar) covalent bonds?

47.  A dipole moment results when a polar molecule has a center for positive charge separate from a center for negative charge

48. Guidelines for Drawing Lewis Structures 1. Determine the total number of valence electrons; (for ions adjust for charge). 2. Arrange atoms in a skeleton structure and connect them with single bonds. 3. Complete octets of the outer atoms 4. If not all of the valence electrons have been used place any extra electrons on the central atom. 5. If the central atom does not have an octet, use lone pairs from terminal atoms to form multiple bonds. -Only C, N, O and sometimes S or P (in combination with C or O) form multiple bonds. Note: If more than one acceptable Lewis structure can be drawn they are called resonance structures.

49. Writing Lewis Structures

50. Section 12.4 Structure of Molecules [NO 2 ] - -

51. B. The VSEPR Model

52. Polar vs. Non-Polar Molecules If a molecule has a center of positive charge that is at a different location than the center of negative charge then the molecule has a dipole and is a polar molecule. H 2 HCl H 2 O

53. Section 11.2 The Hydrogen Atom A. The Energy Levels of Hydrogen Quantized Energy Levels –Since only certain energy changes occur the H atom must contain discrete energy levels.

54. Section 11.3 Atomic Orbitals A. The Hydrogen Orbitals Each principal energy level is divided into sublevels. Hydrogen Energy Levels –Labeled with numbers and letters –Indicate the shape of the orbital 3d3p 4p 3s 4s 2p 4d4f 1s 2s

55. Section 11.3 Atomic Orbitals Sublevels & Orbitals The s sublevel has 1 orbital which is spherically symmetric. The p sublevel has 3 orbitals, shaped like dumbbells (or double tear drops).

56. Section 11.3 Atomic Orbitals Pauli Exclusion Principle - an atomic orbital can hold a maximum of 2 electrons and those 2 electrons must have opposite spins. -Another way to say it is that no 2 electrons can have the same 4 “quantum numbers”: can’t have the same principal energy level (n), same sublevel (s,p,d,f), same orbital (e.g. p x, p y, p z ), AND the same spin (+1/2 or -1/2). B. The Wave Mechanical Model: Further Development

57. Section 11.4 Electron Configurations and Atomic Properties A. Electron Arrangements in the First 18 Atoms on the Periodic Table

58. Section 11.4 Electron Configurations and Atomic Properties Orbital filling and the periodic table B. Electron Configurations and the Periodic Table

59. Section 11.4 Electron Configurations and Atomic Properties A. Naming Compounds That Contain a Metal and a Nonmetal Type 2 Binary Ionic compounds Since the metal ion can have more than one charge, a Roman numeral is used to specify the charge. Ex: gold (I) chloride & gold (III) chloride

60. Section 11.4 Electron Configurations and Atomic Properties B. Naming Binary Compounds That Contain Only Nonmetals Type 3 Compounds