Chapter 1 Elements and Compounds

1. Measurements and observations are made in the macroscopic world. We interpret these measurements and observations using the microscopic world. Chemistry is involved in both worlds. Chapter 1Elements and Compounds

2. Chemistry: A Definition • Find patterns • Develop models • Make predictions • Be quantitative • Experimental testing

3. Elements, Compounds and Mixtures • Elements • Compounds • Mixtures • Homogeneous • Heterogeneous

4. Atomic Symbols • Shorthand notation • Some derived from English names. • Bi for bismuth. • Others derived from non-English names. • Fe for iron. • Upper and lower case important. • CO and Co are different.

5. ChemicalFormula • Atomic symbols and subscripts • Compounds • Molecular • Ionic • Molecular elements • e. g. O2, H2

6. Evidence for the Existence of Atoms • Dalton’s Atomic Theory • Indestructible atoms • Elemental identity • Elemental distinction • Atomic combination

7. The Role of Measurement in Chemistry • Scientific Knowledge • Products • Processes

8. The Role of Measurement in Chemistry • Measurements • Numbers and units • Scientific notation • SI and other units • Prefixes • Conversion factors

9. The Structure of Atoms • Beyond Dalton’s Theory • Three subatomic particles that are important for chemists: • Electron • Proton • Neutron

10. The Structure of Atoms • Absolute vs Relative Charges • Electron charge = Proton charge (with the sign reversed). • Nucleus holds the protons and neutrons.

11. Atomic Number and Mass Number • Both positive integers. • Mass Number (A) Atomic Number (Z) • Z = number of protons. • A = Z + number of neutrons. • X is Atomic Symbol.

12. Isotopes • Same Z • Different A • Identical chemistry

13. Isotopes • Some elements have only one. • Some elements have a few. • Some elements have many. • % natural abundance is the percentage of atoms occurring as a given isotope.

14. Isotopes • Mass of an atom • Absolute mass • Uses a mass unit: gram, ounce, … • Rarely used • Relative mass • Relative to specific isotope of carbon By convention 12C = 12.00000... amu • Ratio called atomic mass. • Used frequently: amu

15. The Difference Between Atoms and Ions • Atoms are neutral. • Ions are charged. • Positively charged ions: cations. • Negatively charged ions: anions.

16. Polyatomic Ions • Many polyatomic anions. • A few polyatomic cations.

17. Polyatomic ions • Spelling • Chemical Formula • Charge Table 1.6

18. The Periodic Table • Atomic number • Groups • Periods Figure 1.8

19. The Periodic Table • Metals • Nonmetals • Semimetals (metalloids) Important to know where an element is situated in the periodic table.

20. The Macroscopic, Atomic and Symbolic Worlds of Chemistry Figure 1.9

21. The Mass of an Atom • Average masses reported. • Average masses are weighted averages. • amu used.

22. Chemical Reactions and the Law of Conservation of Atoms • Fundamental law of chemistry: • Conservation of Mass • What does it mean? • In a chemical reaction, matter is neither created nor destroyed.

23. Chemical Reactions and the Law of Conservation of Atoms • Established empirically. • May not be true, but no counter example has ever been found. • Atomic model based on this law. • Example of scientific method.

24. Chemical Equations as a Representation of Chemical Reactions • Heart of Chemistry • Chemical formulas used. • An arrow is used to separate reactants and products. • Phase information is sometimes included. • Equation carries no implication as to how fast the reaction occurs.

25. Chemical Equations as a Representation of Chemical Reactions Reactants → Products 2Mg(s) + O2(g) → 2MgO(s)

26. Balancing Chemical Equations Notice the equation did not read Mg(s) + O2(g) → MgO(s). Why not?

27. Balancing Chemical Equations Notice the equation did not read Mg(s) + O2(g) → MgO(s). Why not? Chemical equations must be balanced.