Lecture 2

1. Lecture 2 Elements and Atoms B.1-G.4 and 1.1 26-Aug Assigned HW B.4, B.6, B.8, B.10, B.12, B.14, B.18, B.20 Due: Monday 30-Aug

2. A.1-A.3 Concepts • The Scientific Method • Matter (solid, liquid, gas) • Substances (pure matter) • Atoms (simplest matter) • Compound (mixture of elements) • Molecule (smallest unit of a compound) • Units – ALWAYS remember units! • SI Units • Extensive (quantity dependent) and Intensive properties • Error and Sig Figs • Force • Energy • Kinetic, potential, coulombic • Conserved!

3. Origin of Chemists • First chemists were interested in applications • Not so much principles • Alchemists • Blast furnaces, used for extracting iron from iron ore came around 1300 AD. • Physical explanation for combustion marked the entrance to the modern age of chemistry

4. John Dalton • English school teacher • Chemist • Meterologist • Physicist • Color (colour) blind (Daltonism) • Not a good experimentalist

5. Dalton and Atomic Theory • John Dalton used observations of chemical combustion to formulate a hypothesis: • Each element is composed of minute, indivisible particles called atoms. • All atoms of an element are alike in mass and other properties, but different from those of other elements. • Compounds are specific combinations of different elements. • e.g. AB, A2B, AB2 • Atoms are not created or destroyed during a chemical reaction. • AB + AB2 A2B3

6. Discovery of the Electron • The ratio of the mass of an electron to the charge was determined in 1897 by J.J. Thomson Charges separated to anode and cathode. A small perforation allows a beam of cathode rays to pass through and detected on zinc-sulfide screen

7. Discovery of the Electron Cathode rays bent by magnetic fields Cathode Rays soon became known as electrons By varying the field strengths, the m/e ratio was calculated This line of experiments lead Thompson to conclude that the atom was a cloud of (+) and (-) particles.

8. Discover of the Electron • Robert Millikan designed an apparatus that allowed him to observe tiny charged oil droplets (1909) • By varying the electric field necessary to overcome the gravitational force, he calculated the charge of an electron • e = -1.6022 x 10-19 C • Plum pudding model developed by J.J Thompson. Wanna see a video?

9. Discovery of the Nucleus • In 1908, Ernest Rutherford and Hans Geiger used α-particles to study the inner structure of an atom. • Expected most of the particles to pass through, with some scattering modestly. Click here to for a video

10. Discovery of the Nucleus • Observations: Note that this is a perfect example of the scientific method Expected Actual

11. Discovery of the Nucleus

12. Nuclear Model of the Atom • The majority of an atom’s mass is contained in the nucleus • Nucleus surrounded by a ‘cloud’ of electrons • Accounts for most of an atom’s volume • Protons – positively charged particles • Neutrons – charge neutral • Electrons – negatively charged particles • (+) and (-) charges are balanced • Mercury has 80 protons, how many electrons?

13. Protons and Atomic Number • The number of protons in a given atom is called the Atomic Number….denoted by Z ZX • Every atom has a unique Atomic Number! • Any list of the elements can be complimented by atomic number and used for identification purposes. • Most common example is the Periodic Table of the Elements

14. Protons and Atomic Number

15. Protons and Atomic Number • Atomic number of: • Hydrogen (H) • Helium (He) • Carbon (C) • Calcium (Ca) • Krypton (Kr) • Use the atomic number to identify the element: • 29 • 82 • 58

16. Atoms and Mass • Atomic number was first determined unambiguously by Moseley by bombarding elements with rapidly moving electrons and using the properties of the emitted x-rays. • We now have a better way to do it • Mass Spectrometry • Mass of • Hydrogen  1.67 x 10-27 kg • Carbon  1.99 x 10-26 kg

17. Atoms and Mass • Example: How many atoms of lead are in a 25g sample (mass of lead  3.44 x 10-25 kg)?

18. Mass of Sub-atomic Particles

19. Atoms and Mass • We’re looking at a representative mass spectrum of neon (10Ne) • Why is the mass not 10? • Z = 10 • Mass of an electron is insignificant • Neutrons!

20. Isotopes • Isotopes are atoms with the same mass number but a different number of neutrons. • Why aren’t atoms with a different number of protons considered isotopes? • Mass number is the total number of protons and neutrons in a nucleus • Convenient way to determine the number of protons and neutrons

21. Isotopes Textbook Example B.7

22. Isotopes Mass numbers are commonly ignored when writing chemical symbols. Tell me why.

23. Organization of the Elements What’s up with these numbers? One of my favorite links Interactive periodic table

24. Atomic Mass Units • You’ll commonly notice that atomic masses are written as atomic mass units (amu or u) or Daltons (Da) • By international convention, 1 amu is defined as 1/12 the mass of a 12C (Carbon-12). • Mass number is approximately mass in amu.

25. Elemental Organization These react with water

26. Elemental Organization - Conduct electricity, lustrous, malleable and ductile - Some metal properties, but behaves like a non-metal - Does not conduct, neither malleable or ductile