1. Chapter 1: Chemistry and the Atomic/Molecular View of Matter
Chemistry: The Molecular Nature of Matter, 6E
Jespersen/Brady/Hyslop

2. Why Study Chemistry? In every aspect of our modern life
Long life batteries
Materials & miniaturization
Cell phones/pagers
Laptops
Synthetic fibers
Dyes
CDs/DVDs—silicon wafers
Medications
DNA sequencing
Touches all areas of science
Ch 1.1 Chemistry and Its Place among the Sciences

3. Chemistry and the Sciences
Study of matter & its transformations
Seeks answers to fundamental questions about:
What makes up materials that compose our world
How composition affects properties of substances
How substances change when they interact with each other = Chemical Reactions

4. Chemistry and the Sciences
Seeks to understand:
Underlying structures of matter
Forces that determine properties that we observe
Apply this knowledge to:
Create new materials not found in nature
Understand fundamental biological processes

5. Scientific Method
Approach to gathering information & formulating explanations.
Scientists perform experiments in laboratories under controlled conditions
Make observations/collect data
Empirical fact
Something we see, hear, taste, feel, or smell
Something we can measure in laboratory
Organize data so we can see relationships
1.2 Laws and Theories: The Scientific Method

6. Scientific Method Law or Scientific Law Hypothesis
Broad generalization
Based on results of many experiments
Only states what happens
Doesn’t explain why they happen
Hypothesis
Mental picture that explains observed laws
Tentative explanation of data
Make predictions
Devise experiments to test
Go back to laboratory & perform

7. Scientific Method Theory Tested explanation of how nature behaves
Devise further tests
Depending on results, may have to modify theory
Can never prove theory is absolutely correct

8. Scientific Method Ex. Study gases
Discover Volume (V) of gas depends on
Pressure (P)
Temperature (T)
Amount (n)
Data
Recorded observations of relationship between V, P, T & n
Law
R = constant
Kinetic Theory of Gases
Explains gas behavior (Ch 11)

9. Atomic Theory Most significant theoretical model of nature Atoms
Tiny submicroscopic particles
Make up all chemical substances
Make up everything in Macroscopic world
Smallest particle that has all properties of given element
Composed of:
Electrons
Neutrons
Protons

10. Matter & Its Classifications
Anything that has mass & occupies space
Mass
How much matter given object has
Measure of object’s momentum, or resistance to change in motion
Weight
Force with which object is attracted by gravity
Ex. Mass vs. Weight
Astronaut on moon & on earth
Weight on moon = 1/6 weight on earth
Same mass regardless of location

11. Matter Chemical Reactions Decomposition Ex.
Transformations that alter chemical compositions of substances
Decomposition
Chemical reaction where 1 substance broken down into 2 or more simpler substances
Ex.
Electric current
Sodium metal + chlorine gas
Molten sodium chloride

12. Elements
Substances that can’t be decomposed into simpler materials by chemical reactions
Substances composed of only 1 type of atom
Simplest forms of matter that we can work with directly
More complex substances composed of elements in various combinations
diamond = carbon
gold
sulfur

13. Chemical Symbols for Elements
One or two letter symbol for each element name
First letter capitalized, second letter lower case
Ex. C = carbon S = sulfur
Ca = calcium Ar = argon
Br = bromine H = hydrogen
Cl = chlorine O = oxygen
Used to represent elements in chemical formulas
Ex. Water = H2O
Carbon dioxide = CO2
Most based on English name
Some based on Latin or German names

14. Chemical Symbols English Name Chemical Symbol Latin Name Sodium Na
Natrium
Potassium K
Kalium
Iron Fe
Ferrum
Copper Cu
Cuprum
Silver Ag
Argentum
Gold Au
Aurum
Mercury Hg
Hydrargyrum
Antimony Sb
Stibium
Tin Sn
Stannium
Lead Pb
Plumbum
Tungsten W
Wolfram (German)

15. Compound Formed from 2 or more atoms of different elements
Always combined in same fixed ratios by mass
Can be broken down into elements by some chemical changes
Ex. Water decomposed to elemental hydrogen & oxygen
Mass of oxygen = × mass of hydrogen

16. Pure Substance vs. Mixture
Pure substances
Elements and compounds
Composition always same regardless of source
Mixture
Can have variable compositions
Made up of two or more substances
Ex. CO2 in water—varying amounts of “fizz” in soda
2 broad categories of mixtures:
Heterogeneous
Homogeneous

17. Homogeneous Mixtures Same properties throughout sample Solution
Thoroughly stirred homogeneous mixture
Ex.
Liquid solution
Sugar in water
Gas solution
Air
Contains nitrogen, oxygen, carbon dioxide & other gases
Solid solution
US 5¢ coin – Metal Alloy
Contains copper & nickel metals

18. Heterogeneous Mixtures
2 or more regions of different properties
Solution with multiple phases
Separate layers
Ex.
Salad dressing
Oil & vinegar
Ice & water
Same composition
2 different physical states

19. Physical Change No new substances formed
Substance may change state or the proportions
Ex. Ice melting
Sugar or salt dissolving
Stirring iron filings & sulfur together

20. Chemical Change or Chemical Reaction Ex.
Formation of new substance or compound
Involves changing chemical makeup of substances
New substance has different physical properties
Can’t be separated by physical means
Ex.
Fool’s gold
Compound containing sulfur & iron
No longer has same physical properties of free elements
Can’t be separated using magnet

21. Learning Check:
For each of the following, determine if it represents a Chemical or Physical Change:
Chemical
Physical
Magnesium burns when heated
Magnesium metal tarnishes in air
Magnesium metal melts at 922 K
Grape Kool-aid lightens when water is added X X X X

22. Classification of Matter

23. Learning Check: Classification
Hot
Cocoa
Ice
(H2O)
White Flour
Table Salt (NaCl)
Pure substance
Element
Compound
Molecule
Heterogeneous Mixture
Homogeneous Mixture X X X X X X X

24. Law of Definite Proportions
In given compound, elements always combine in same proportions by mass.
Ratio of masses of each element is fixed for given compound
Implication:
Each atom has fixed specific mass
Ex. Fool’s gold, pyrite, iron (III) sulfide
Mass ratio always
1.00 g of Iron to g of Sulfur
Ex. Water
Mass ratio always: 8 g O to 1 g H
1.4 Dalton and the Atomic Theory

25. Law of Conservation of Mass
No detectable gain or loss of mass occurs in chemical reactions.
Mass is conserved.
Implication:
Reactions involve reorganization of materials.
Mass before & after reaction is unchanged.
For reaction in sealed container, mass of reactants = mass of products

26. Learning Check: Chemical Laws
Magnesium burns in oxygen to form magnesium oxide. If g of Mg are consumed and g of MgO are produced, what mass of oxygen was consumed?
28.00 g – g = 11.12g O

27. Learning Check: Chemical Laws
In a sample of MgO, there are g Mg and g O. What mass of O would there be in a sample that contains 2.00 g of Mg?
X = 1.30 g O

28. Dalton’s Atomic Theory
John Dalton
Developed underlying theory to explain
Law of Conservation of Mass
Law of Definite Proportions
Reasoned that if atoms exist, they have certain properties
Dalton’s Atomic Theory
Matter consists of tiny particles called atoms.

29. Dalton’s Atomic Theory (cont)
Atoms are indestructible.
In chemical reactions, atoms rearrange but do not break apart.
In any sample of a pure element, all atoms are identical in mass & other properties.
Atoms of different elements differ in mass & other properties.
In given compound, constituent atoms are always present in same fixed numerical ratio.

30. Proof Of Atoms Early 1980’s, use Scanning Tunneling Microscope (STM)
Surface can be scanned for topographical information
Image for all matter shows spherical regions of matter
Atoms
Fig 1.10
STM of palladium

31. How Do We Visualize Atoms?
Atoms represented by spheres
Different atoms have different colors
Standard scheme given in Fig is represented on the right.
Fig 1.11 in 3 pieces so fits on slide

32. Molecules Atoms combine to form more complex substances
Discrete particles
Each composed of 2 or more atoms
Ex.
Molecular oxygen, O2
Carbon dioxide, CO2
Ammonia, NH3
Sucrose, C12H22O11

33. Chemical Formulas Specify composition of substance Chemical symbols
Represent atoms of elements present
Subscripts
Given after chemical symbol
Represents relative numbers of each type of atom
Ex.
Fe2O3 : iron & oxygen in 2:3 ratio

34. Chemical Formulas Free Elements Ex. Iron Fe Neon Ne Diatomic Molecule
Element not combined with another in compounds
Just use chemical symbol to represent
Ex. Iron Fe Neon Ne
Sodium Na Aluminum Al
Diatomic Molecule
Molecules composed of 2 atoms each
Many elements found in nature
Ex. Oxygen O2 Nitrogen N2
Hydrogen H2 Chlorine Cl2

35. Depicting Molecules Want to show: Three ways of visualizing molecules:
Order in which atoms are attached to each other
3-dimensional shape of molecule
Three ways of visualizing molecules:
Structural formula
Ball-and-Stick model
Space filling model

36. 1. Structural Formulas Use to show how atoms are attached
Atoms represented by chemical symbols
Chemical bonds attaching atoms indicated by lines
H2O
water
CH4
methane

37. 3-D Representations of Molecules
Hydrogen molecule, H2
Oxygen molecule, O2
Nitrogen molecule N2
Chlorine molecule,
Cl2
Use touching spheres to indicate molecules
Different colors indicate different elements
Relative size of spheres reflects differing sizes of atoms
Fig 1.12

38. 2. “Ball-and-Stick” Model
Spheres = atoms
Sticks = bonds
Chloroform, CHCl3
Methane, CH4

39. 3. “Space-Filling” Model
Shows relative sizes of atoms
Shows how atoms take up space in molecule
Methane CH4
Water H2O
Chloroform, CHCl3

40. More Complicated Molecules
Sometimes formulas contain parentheses
How do we translate into a structure?
Ex. Urea, CO(NH2)2
Expands to CON2H4
Atoms in parentheses appear twice
Ball-and-stick model
Space-filling model

41. Hydrates Crystals that contain water molecules Dehydration
Ex. plaster: CaSO4∙2H2O calcium sulfate dihydrate
Water is not tightly held
Dehydration
Removal of water by heating
Remaining solid is anhydrous (without water)
Blue = CuSO4 •5H2O
White = CuSO4

42. Counting Atoms
Subscript following chemical symbol indicates how many of that element are part of the formula
No subscript implies a subscript of 1.
Quantity in parentheses is repeated a number of times equal to the subscript that follows.
Raised dot in formula indicates that the substance is a hydrate
Number preceding H2O specifies how many water molecules are present.

43. Counting Atoms Ex. 1 (CH3)3COH Subscript 3 means 3 CH3 groups
So from(CH3)3, we get 3 × 1C = 3C
3 × 3H = 9H
#C = 3C + 1C = 4 C
#H = 9H + 1H = 10 H
#O = 1 O
Total # of atoms = 15 atoms

44. Counting Atoms Ex. 2 CoCl2 · 6H2O
The dot 6H2O means you multiple both H2 & O by 6
So there are:
#H 6 × 2 = 12 H
#O 6 × 1 = 6 O
#Co 1 × 1 = 1 Co
#Cl 2 × 1 = 2 Cl
Total # of atoms = 21 atoms

45. Your Turn!
Count the number of each type of atom in the chemical formula given below
Na2CO3
(NH4)2SO4
Mg3(PO4)2
CuSO4∙5H2O
(C2H5)2N2H2
___Na, ___ C, ___ O
___N, ___H, ___S, ___O
___Mg, ___P, ___O
___Cu, ___S, ___O, ___H
___C, ___H, ___N 2 1 3 2 8 1 4 3 2 8 1 1 9 10
Dialog: How do I count the atoms of elements in a chemical formula? 4 12 2

46. Dalton’s Atomic Theory
We now have the tools to explain this theory & its consequences
All molecules of compound are alike & contain atoms in same numerical ratio.
Ex. Water, H2O
Ratio of oxygen to hydrogen is 1 : 2
1 O atom : 2 H atoms in each molecule
O weighs 16 times as much as H
1 H = 1 mass unit
1 O = 16 mass units

47. Atoms in Fixed Ratios by Mass
For water in general:
mass O = 8 mass H
Regardless of amount of water present

48. Dalton’s Atomic Theory
Successes:
Explains Law of Conservation of Mass
Chemical reactions correspond to rearranging atoms.
Explains Law of Definite Proportions
Given compound always has atoms of same elements in same ratios.
Predicted Law of Multiple Proportions
Not yet discovered
Some elements combine to give 2 or more compounds
Ex. SO2 & SO3

49. Law Of Multiple Proportions
When 2 elements form more than one compound, different masses of one element that combine with same mass of other element are always in ratio of small whole numbers.
Atoms react as complete (whole) particles.
Chemical formulas
Indicate whole numbers of atoms
Not fractions
Dialog: What is the law of multiple proportions?

50. Using Law Of Multiple Proportions
sulfur sulfur
dioxide trioxide
Mass S g g
Mass O g g
Use this data to prove law of multiple proportions
Each molecule has one sulfur atom, and therefore the same mass of sulfur. The oxygen ratio is 3 to 2 by both mass and atoms

51. Law of Multiple Proportions
Compound
Sample Size
Mass of Sulfur
Mass of Oxygen
Sulfur dioxide
64.06 g
32.06 g
32.06 g
Sulfur trioxide
80.06 g
32.06 g
48.00 g
Ratio of

52. Molecules Small and Large
So far we’ve only discussed small molecules
Some are very large, especially those found in nature
Same principles apply to all
Ex. DNA - short segment

53. How Do We Know Formulas? Hardly “out of the blue”
Don’t know formula when compound 1st isolated
Formulas & structures backed by extensive experimentation
Use results of experiments to determine
Formula
Chemical reactivity
Molecular Shape
Can speculate once formula is known
Determine from more experiments

54. Visualizing Mixtures Look at mixtures at atomic/molecular level
Different color spheres stand for 2 substances
Homogeneous mixture/solution – uniform mixing
Heterogeneous mixture – 2 phases a. b.
Fig. 1.21

55. Chemical Reactions
When 1 or more substances react to form 1 or more new substances
Ex. Reaction of methane, CH4, with oxygen, O2, to form carbon dioxide, CO2, & water, H2O.
Reactants = CH4 & O2
Products = CO2 & H2O
How to depict?
Words too long
Pictures too awkward
Fig. 1.23

56. Chemical Equations
Use chemical symbols & formulas to represent reactants & products.
Reactants on left hand side
Products on right hand side
Arrow () means “reacts to yield”
Ex. CH4 + 2O2  CO2 + 2H2O
Coefficients
Numbers in front of formulas
Indicate how many of each type of molecule reacted or formed
Equation reads “methane & oxygen react to yield carbon dioxide & water”

57. Conservation of Mass in Reactions
Mass can neither be created nor destroyed
This means that there are the same number of each type of atom in reactants & in products of reaction
If # of atoms same, then mass also same
CH O  CO H2O
4 H + 4O + C = H + 4O + C

58. Balanced Chemical Equation
Ex C4H O2  8CO H2O
1 C & 2 O per molecule
2 H & 1 O per molecule
4 C & 10 H per molecule
2 O per molecule
Subscripts
Define identity of substances
Must not change when equation is balanced

59. Balanced Chemical Equation
Ex. 2C4H O2  8CO H2O
13 molecules of O2
8 molecules of CO2
10 molecules of C4H10
2 molecules of C4H10
Coefficients
Number in front of formulas
Indicate number of molecules of each type
Adjusted so # of each type of atom is same on both sides of arrow
Can change

60. Balanced Chemical Equations
How do you determine if an equation is balanced?
Count atoms
Same number of each type on both sides of equation?
If yes, then balanced
If no, then unbalanced
Ex. 2C4H O2  8CO H2O
Reactants Products
2×4 = 8 C 8×1 = 8 C
2×10 = 20 H 10×2 = 20 H
13×2 = 26 O (8×2)+(10×1)= 26 O

61. Learning Check Fe(OH)3 + 2 HNO3  Fe(NO3)3 + 2 H2O Not Balanced
Only Fe has same number of atoms on either side of arrow.
Reactants
Products Fe 1
3 + (2×3) = 9
(3×3) + 2 = 11 O
3 + 2 = 5
(2×2) = 4 H 2 3 N

62. Your Turn!
How many atoms of each element appear on each side of the arrow in the following equation? 4NH3 + 3O2 → 2N2 + 6H2O
Reactants
Products N
(4 × 1) = 4
(2 × 2) = 4 O
(3 × 2) = 6
(6 × 1) = 6 H
(4 × 3) = 12
(6 × 2) = 12

63. Your Turn!
Count the number of atoms of each element on both sides of the arrow to determine whether the following equation is balanced.
2(NH4)3PO4 + 3Ba(C2H3O2)2 → Ba3(PO4)2 + 6NH4C2H3O2
Reactants
Products N
(2 × 3) = 6
(6 × 1) = 6 H
(2×3×4)+(3×3×2) = 42
(6×4) + (6×3) = 42 O
(2×4) + (3×2×2) = 20
(2×4) + (6×2) = 20 P
(2 × 1) = 2 Ba
(3 × 1) = 3