Atomic Structure and Classification of Matter

1. ATOMIC STRUCTURE Chapter 4

2. Notes sheet 1

3. What is matter? Matter is defined as anything that has mass and volume. Atoms are defined as the basic units of matter. NOTE: Energy is NOT matter and does not have mass or volume. Energy and matter are related by E=mc2

4. matter pure substances mixtures homogeneous mixtures heterogeneous mixtures elements compounds uniform (solution) non-uniform (separate parts) one type of atom different atoms bonded to each other Hg element I2 element HgI2 compound

6. Classifying Matter H2O

7. Compounds vs. mixtures • Compounds are groups of atoms that are held together by chemical bonds. A molecule is the smallest unit of a compound. • The elements within a compound can only be separated by a chemical change, which breaks the bonds in the compound. • Mixtures, however, can be separated by physical methods such as filtration or evaporation, which DO NOT chemically alter the substances.

8. Compounds and Diatomics • Compounds are groups of two or more different atoms that are held together by chemical bonds. • There are some gases that are bound to themselves as pure elements (diatomic molecules): Hydrogen = H2 Nitrogen = N2 Oxygen = O2 Fluorine = F2 Chlorine = Cl2 Bromine = Br2 Iodine = I2

9. How would you classify these?

10. Chemical vs. Physical Properties

11. Chemical vs. Physical Changes • Properties of matter are characteristics that can be tested or observed and are used to identify matter. • Physical properties are those which can be observed WITHOUT changing the chemical make-up of the matter. They are observed during physical changes. • Chemical properties can only be observed when matter is involved in a chemical reaction, which CHANGES the chemical composition. Chemical changes will create a new substance whereas physical changes do not.

12. Chemical vs. Physical Changes

13. Chemical vs. Physical Changes • In a physical change, the chemical identity of the substance ___________________. • In a chemical change, the chemical identity of a substance __________. • Another name for a chemical change is a ______________ _________________. doesn’t change changes chemical reaction

14. More Examples of Properties • Carbon reacts with oxygen to form carbon dioxide and water. • Water boils at 100ºC. • Paper burns. • A solution of KNO3 is colorless. • Iron rusts. • Solid sulfur is dull and yellow. • Gold has a very high density (19.3 g/cm3).

15. Sheet 2

16. Matter is made of atomos of air, fire, wind and earth! 460 BC Dalton 1803 Thomson 1897 Rutherford 1909 Bohr 1913 Present Day

18. The Law of Conservation of Mass Antoine Lavoisier (1700s) compared the masses of substances before and after a reaction and found that the mass always remained constant. This led to the Law of Conservation of Mass. In other words, matter cannot be created or destroyed, only rearranged in chemical reactions.

19. The Law of Definite Proportions The same samples of a pure compound always contain elements in the same mass proportion. French chemist Joseph Proust

20. Dalton’s Atomic Theory The idea of an indivisible thing that made up all matter was refined by chemist John Dalton in 1803. His ideas were:

21. Daltons Law of Multiple Proportions: • Multiple Proportions: Since atoms bond in small, whole number ratios to form compounds, their masses are small whole number ratios. Ex: CO vs. CO2

22. Thomson’s cathode ray tube and plum pudding model electron beam

23. Ernest Rutherford • While studying radioactive elements, New Zealander Physicist Ernest Rutherford found that radioactive alpha particles deflected when fired at a very thin gold foil. • This was known as the gold foil experiment, and it suggested that the atom was not a hard sphere as thought, but was mostly space, with a small concentration of mass. • This concentration of mass became known as the nucleus. • Link to experiment…

24. Millikan’s oil-drop experiment • Robert Millikan’s classic oil-drop experiment allowed the charge of a single electron to be determined: 1.60 x 10-19 C. • Using these two numbers, we can calculate the mass of an electron: The mass of an electron is about 1/2000 of the mass of a proton!

25. Rutherford’s gold foil exp. and new atomic model

26. The Bohr Model Neils Bohr, a Russian scientist, proposed that the electrons must be in specific energy “orbits” (called electron shells) around the nucleus. Lower energy shells closer to the nucleus

27. The Modern Electron Cloud (aka wave-mechanical) Model Democritus and Dalton’s atom Thompson’s electrons Rutherford’s space and nucleus Bohr’s energy levels Chadwick’s neutrons (not to scale) Most atoms ~ 1-5 Å (Angstroms) = (1x10-10m)

28. Sheet 3

29. Elements O 15.9994 • We currently know of about 110 elements, 92 of which are naturally occurring. • We illustrate an element with an atomic symbol. • The atomic number tells the number of protons and identifies the element. • The mass number is the total number of the protons plus the neutrons. (on the P. T. the average atomic mass of isotopes is given) 8 OXYGEN atomic number average atomic mass

30. Isotope symbols a charge is shown here for an ion. This tells the number of electrons that have been gained or lost +2 • Isotope symbols give the element symbol, the mass and atomic numbers. From these, the number of protons, neutrons, and electrons can be determined.

31. What element is this? What is the mass of this isotope? hey, does this proton make my MASS look big?

32. Ions Cl Cl- • An ion is an atom that has gained or lost one or more electrons. • When an atom bonds with another atom, it seeks to gain electrons or lose them. For instance: • Cl has 7 and will gain one electron • Na has 1 and will lose one electron Na Na+ Positive ions are called “cations” Negative ions are called “anions”

34. Practice • How many protons, neutrons, and electrons are present in (a) 27Al3+ (b) 79Se2- • Write the isotope notation for an ion that contains 20 protons, 21 neutrons, and 18 electrons. • Write the isotope notation for an atom of lead that has 128 neutrons. • Write the isotope notation for the following:

35. Formula mass (aka molar mass) The formula mass for a compound is the sum of the ______ ______of all elements in the formula. Ex: Calculate the formula mass for Al2(SO4)3 atomic masses

36. Notes 4

37. Weighted average mass 91.0% 25.2 kg 5.0% 16.1 kg 3.0% 10.0 kg ? % 12.0 kg

38. Natural Abundance - Isotopes • There is not just one type of each atom, there are several. When an atom has more or less neutrons than another atom of the same element, we call them isotopes. • For instance, the element carbon has 6 protons, but it could have 5, 6, 7, or 8 neutrons, to form Carbon-11, Carbon-12, Carbon-13, and Carbon-14. Each has a different mass. • In nature, there is a mix of different natural isotopes. We use this mix to calculate average atomic mass. ex: carbon is 12.011 amu 12C 13C 14C

39. Calculating Average Atomic Mass • To find average atomic mass, we multiply the relative abundance of an isotope by the mass of the isotope. We then add each of the products for each isotope.

40. Average atomic mass for one mole of atoms

41. Isotopes, Ions, and Allotropes (Oh my) • Isotopes are atoms of the same element with different numbers of neutrons. • Ions are atoms of the same element with different numbers of electrons. (Ions are easy to create by adding or removing electrons from a neutral atom). • Allotropes are forms of the same element, but bonded in different structures. • Diamond and pencil graphite are examples of allotropes. They are both pure carbon, but in different structures.

42. Isotopes and allotropes

43. Sheet 5

44. Nuclear Chemistry and the Band of Stability • What particles make up the nucleus? • What is the charge and function of each particle? #n > #p protons and neutrons #n = #p proton (+) -determines the identity of the atom neutron (0) –helps hold nucleus together

45. Elements above the band of stability need to decrease their n/p ratio Elements below the band of stability need to increase their n/p ratio

46. Why do nuclear reactions occur? • Nuclear reactions occur when a nucleus becomes unstable. • Protons and neutrons are attracted to each other by the strong nuclear force. In a stable nucleus, the attraction due to the strong force is greater than the repulsion due to electrostatic force. As elements get heavier, they become more unstable. Extra neutrons must be present to the nucleus (like glue) to increase stability by increasing the strong force. • Nuclear reactions are DIFFERENT from chemical reactions, because new elements form.

47. What is radioactivity? • Radioactivity is the spontaneous emission of radiation from an element to achieve a more stable state. • Uranium was the first radioactive element isolated (by Bequerel), followed by radium and polonium (by Marie Curie and her husband Pierre). • There are no stable isotopes for elements after Bismuth (#83).

48. Types of Radiation

50. Nuclear Reactions • Nuclear reactions change the composition of an atom’s nucleus –the element will change!! • Examples of naturally occurring nuclear reactions include alpha and beta decay, and fission and fusion. • Some nuclei can become unstable by artificial transmutation, where a nucleus is bombarded (or shot) with a particle that creates instability and causes radioactive decay. • Nuclear reactions can produce enormous amounts of energy as nuclear mass is converted into energy (E=mc2)