1. Figure: 03-02
Title:
The difference between a subscript in a chemical formula and a coefficient in front of the formula.
Caption:
Notice how adding the coefficient 2 in front of the formula (line 2) has a different effect on the implied composition than adding the subscript 2 to the formula (in line 3). The number of atoms of each type (listed under composition) is obtained by multiplying the coefficient and the subscript associated with each element in the formula.

2. Figure: 03-03
Title:
Methane reacts with oxygen to produce the flame in a Bunsen burner.
Caption:
The methane (CH4) in natural gas and oxygen (O2) from the air are the reactants in the reaction, while carbon dioxide (CO2) and water (H2O) are the products.

3. Figure: 03-04
Title:
Balanced chemical equation for the combustion of CH4.
Caption:
The drawings of the molecules involved call attention to the conservation of atoms through the reaction.

4. Figure: UNSE3.1
Title:
Sample Exercise 3.1
Caption:
Interpreting and balancing chemical equations.

5. Figure: UNPE3.1
Title:
Practice Exercise 1
Caption:
Law of conservation of mass.

6. Figure: 03-05
Title:
Combustion of magnesium metal in air.
Caption:
When magnesium metal burns, the Mg atoms react with O2 molecules from the air to form magnesium oxide, MgO, an ionic solid: 2 Mg(s) + O2(g) → 2MgO(s). The photos show what we see in the laboratory. The ribbon of magnesium metal (left) is surrounded by oxygen in the air, and as it burns an intense flame is produced. At the end of the reaction, a rather fragile ribbon of white solid, MgO, remains. The models show the atomic-level view of the reactants and products.

7. Figure: UNT01
Title:
Table 3.1
Caption:
Combination and Decomposition Reactions

8. Figure: 03-08
Title:
Comparing the mass of 1 molecule H2O and 1 mol H2O
Caption:
Notice that the masses are numerically equal but have different units (18.0 amu compared to 18.0 g) representing the huge difference in mass.

9. Figure: UNT02
Title:
Table 3.2
Caption:
Mole Relationships

10. Figure: 03-10
Title:
Procedure for interconverting the mass and the number of formula units of a substance.
Caption:
The number of moles of the substance is central to the calculation; thus, the mole concept can be thought of as the bridge between the mass of a substance in grams and the number of formula units.

11. Figure: 03-11
Title:
Procedure for calculating an empirical formula from percentage composition.
Caption:
The central part of the calculation is determining the number of moles of each element in the compound. The procedure is also summarized as “percent to mass, mass to mole, divide by small, multiply ‘til whole.”

12. Figure: 03-12
Title:
Apparatus to determine percentages of carbon and hydrogen in a compound.
Caption:
The compound is combusted to form CO2 and H2O. Copper oxide helps to oxidize traces of carbon and carbon monoxide to carbon dioxide and to oxidize hydrogen to water.

13. Figure: UNT03
Title:
Table 3.3
Caption:
Information from a Balanced Equation

14. Figure: UN
Title:
Grams reactant to grams product conversion sequence.
Caption:
Flow chart showing the sequence of calculations used to find the number of grams of product for the combustion of butane.

15. Figure: 03-13
Title:
The procedure for calculating amounts of reactants or products in a reaction.
Caption:
The number of grams of a reactant consumed or of a product formed in a reaction can be calculated, starting with the number of grams of one of the other reactants or products. Notice how molar masses and the coefficients in the balanced equation are used.

16. Figure: UN
Title:
Sample Exercise 3.16
Caption:
Calculating amounts of reactants and products.

17. Figure: 03-14
Title:
Increasing concentration of atmospheric CO2.
Caption:
The worldwide concentration of CO2 has increased from about 290 ppm to over 370 ppm over the past 150 years. The concentration in ppm is the number of molecules of CO2 per million (106) molecules of air. Data before 1958 came from analyses of air trapped in bubbles of glacial ice.

18. Figure: 03-15
Title:
Example illustrating a limiting reactant.
Caption:
Because the H2 is completely consumed, it is the limiting reagent in this case. Because there is a stoichiometric excess of O2, some is left over at the end of the reaction. The amount of H2O formed is related directly to the amount of H2 consumed.

19. Figure: UNE03-01
Title:
Exercise 3.1
Caption:
The reaction between reactant A and reactant B.

20. Figure: UNE03-02
Title:
Exercise 3.2
Caption:
Eight H2 molecules.

21. Figure: UNE03-25a
Title:
Exercise 3.25(a)
Caption:
Benzaldehyde.

22. Figure: UNE03-25b
Title:
Exercise 3.25(b)
Caption:
Vanillin.

23. Figure: UNE03-25c
Title:
Exercise 3.25(c)
Caption:
Isopentyl acetate.

24. Figure: UNE03-03
Title:
Exercise 3.3
Caption:
Products of a decomposition reaction.

25. Figure: UNE03-04
Title:
Exercise 3.4
Caption:
Products of a combustion reaction.

26. Figure: UNE03-26
Title:
Exercise 3.26
Caption:
Calculating the percentage of carbon by mass.

27. Figure: UNE03-06
Title:
Exercise 3.6
Caption:
High-temperature reaction between CH4 and H2O.

28. Figure: UNE03-07
Title:
Exercise 3.7
Caption:
Diagram of a mixture of nitrogen and hydrogen.

29. Figure: UNE03-08
Title:
Exercise 3.8
Caption:
Molecules of NO and O2.

30. Figure: UNE03-05
Title:
Exercise 3.5
Caption:
Glycine.