1. HP Chemistry Slides from DOSAQAs with notes Fall 2014-2015

2. Review graphing practice A line of best fit is a straight (or curved) that best represents the data on a scatter plot. This line may pass through some of the points, none of the points, or all of the points.

3. Shadow (meters) depend ent variable Student Height (meters) independent variable If two quantities are directly proportional, then their ratio is constant. Direct  X/Y = k; i.e. X = kY

4. Volume (ml) dependent variable Pressure (atm) - independent variable If two quantities are inversely proportional, then their product is constant. XY=k; which can also be written as Y=k/X

5. Physical properties PropertyDescriptionExamples Electrical conductivity Ability to carry electricityCopper wiring Heat conductivityAbility to transfer energy as heatAluminum pots and pans DensityMass-to-volume ratio of a substance (how tightly packed a substance is) Lead sinkers for fishing Melting pointTemperature at which solid changes to liquid Ice  liquid water Boiling pointTemperature at which liquid changes to gas Liquid water  vapor MalleabilityAbility to be hammered or beaten into thing sheets Silver jewelry DuctilityAbility to be drawn into a thin wireTantalum dental tools

6. Classification of matter www.chemistrytutors.net/wp-content/uploads/20...

7. Temperature scales and conversions Fahrenheit Celsius (SI system) ° F = 1.8(°C)+ 32 ° C = (5/9)(°F-32) Kelvin (Absolute) Temperature scale K = ° C + 273

8. Celsius Temperature Scale The Celsius temperature scale “ centigrade " scale ("consisting of or divided into 100 degrees.“) Andres Celsius (Swedish;1701-1744) developed the centigrade scale for scientific purposes. 100 degrees between freezing point (0˚C) & boiling point (100˚C) of pure water at sea level air pressure. An international conference on weights and measures voted to name the centigrade scale after its inventor in 1948.

9. The Kelvin scale Based on the Celsius scale, but has no negative numbers. Zero on the Kelvin scale is considered to be absolute zero ; that is, the point at which all-molecular motion stops. K = °C + 273

11.  Law of conservation of mass : mass is neither created nor destroyed during ordinary chemical reactions or physical changes.  Law of definite proportions : a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound.  Law of multiple proportions : if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers.

12. Dalton’s Atomic Theory (1808)  All matter is composed of extremely small particles called atoms.  Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties.  Atoms cannot be subdivided, created, or destroyed.  Atoms of different elements combine in simple whole-number ratios to form chemical compounds.  In chemical reactions, atoms are combined, separated, or rearranged.

13. ParticleSymbolRelative electric charge Mass number Relative mass (amu) Actual mass (kg) Electron e -, 0 -1 e 0 0.00054869.109 x 10 -31 Proton p+, 11Hp+, 11H +1 1 1.007276 1.673 x 10 -27 Neutron n 0, 1 0 n 0 1 1.008665 1.675 x 10 -27 *1 amu (atomic mass unit) = 1.660540 x 10 -27 kg Similar chart on page 76 in your textbook

14.  Atomic number (Z) = the number of protons in the nucleus of an atom of that element.  Z is from the German Zahl or “number.”  Mass number (i.e. atomic mass) [A] = total number of protons and neutrons in the nucleus of an isotope.  A from German Atomgewichte or “atomic weight.”

15. What makes the atomic mass different?  Isotopes = atoms of the same element that have different masses (i.e. different numbers of neutrons). They may be naturally occurring or man-made in a laboratory.  Nuclide is a general term for any isotope of an element

16. Ion: an atom or group of bonded atoms that has a positive or negative charge as a result of having either more or less electrons than the neutral atom from which it originated. Any process that results in the formation of an ion is called ionization. Oxidation number/state/valence: a number assigned to an atom in a molecular compound or molecular ion that indicates the general distribution of electrons among the bonded atoms; i.e. the atom or ion’s charge.

17. Atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any nuclide is determined by comparing it with the mass of the carbon-12 atom. Average atomic mass is the weighted average of the atomic masses of the naturally occurring isotopes of an element. This is what is reported on a periodic table.

18. Some modifications to the theory  Atoms are divisible into even smaller particles (like what?)  Elements can have atoms with different masses (isotopes)

19. Calculating average atomic mass Multiplying the atomic mass of each isotope by its relative abundance (expressed in decimal form) and adding the results: Example: Copper, Cu (0.6917 x 62.929599 amu) + (0.3083 x 64.927793 amu) = 63.55 amu Atomic mass unit, or 1 amu, is exactly 1/12 the mass of a carbon-12 atom. The atomic mass of any nuclide is determined by comparing it with the mass of the carbon- 12 atom. Relative abundance is determined using a mass spectrometer

20. Mass spectrometer mn.wikipedia.org/wiki/ Хэрэглэгч :Bilg... mn.wikipedia.org/wiki/ Хэрэглэгч :Bilg...

21. Accelerator mass spectrometer at Lawrence Livermore National Laboratory

22. ueip.org/mass-spectrometer-isotope-analysis-u...

23. Mass spectrum

25. Discovery of the electron 1. An object placed between the cathode and the opposite end of the tube cast a shadow on the glass. 2. A paddle wheel placed on rails between the electrodes rolled along the rails from the cathode toward the anode. 3. Cathode rays were deflected by a magnetic field in the same manner as a wire carrying electric current, which was known to have negative charge. 4. The rays were deflected away from a negatively charged object. Cathode ray tube

26. These observations led to the hypothesis that the particles that compose cathode rays are negatively charged. This was supported by experiments undertaken by J.J. Thompson in 1897. In one experiment he found that the ratio of the charge of cathode-ray particles to their mass was always the same, regardless of the cathode or the type of gas used. He concluded that cathode rays are composed of identical negatively charged particles that were later named electrons.

27. Millikan’s Oil Drop Experiment (1909) He determined the mass of an electron. Mass of an electron: 9.109 x 10 -31 kg.

28. Discovery of the atomic nucleus Ernest Rutherford, Hans Geiger and Ernest Marsden (1911) They bombarded a very thin piece of gold foil with alpha particles (about four times the mass of a hydrogen atom). They found that the majority of the particles went straight through the foil and a very few were deflected.

29. Therefore, each atom in the gold foil had a very small, dense, positively charged nucleus surrounded by electrons.

30. Discovery of the neutron library.thinkquest.org/27954/neutron.html library.thinkquest.org/27954/neutron.html In 1932, James Chadwick bombarded beryllium (Be) with alpha particles. He allowed the radiation emitted by beryllium to incident on a paraffin wax. It was found that protons were shot out form the paraffin wax. People began to look for what was in the "beryllium radiations".

31. Some people suggested that the radiations may be gamma radiation. However, Chadwick found that the radiation could not be gamma radiation since energy and momentum were not conserved in its production. He showed that all the observations could be explained if the radiation consisted of neutral particles of mass approximately equal to that of proton. This neutral particle was named neutron. Equation of the nuclear reaction

32. This new tool in atomic disintegration did not have to overcome any electric barrier and was capable of penetrating and splitting the nucleus of even the heaviest elements. Chadwick thus prepared the way towards the fission of uranium 235 and towards the creation of the atomic bomb. For this epoch-making discovery he was awarded the Hughes Medal of the Royal Society in 1932, and subsequently the Nobel Prize for Physics in 1935.

33. Nuclear forces Usually like charges repel, but protons and neutrons can be very close together. These short-range proton neutron, proton proton, and neutron neutron forces hold the nuclear particles together And are called nuclear forces.

34. Discovery of the neutron led to the beginning of current theories of nuclear structure. Immediately, the neutron-proton model ( the Rutherford-Bohr model) of the nucleus was adopted:  The nucleus is made up of protons and neutrons. These are bound together by a strong nuclear force.  Electrons and protons carry equal but opposite charges. In a neutral atom, the number of electrons is the same as the number of protons.  Electrons orbit the nucleus at certain fixed levels called shells.

35. The Mole Simply represents a number. Just as the term dozen refers to the number twelve, the mole represents the number 6.02 x 10 23

36. The history of the MOLE Amadeo Avogadro (1811) proposed that equal volumes of different gases at the same temperature contain equal numbers of molecules.

37. Stanislao Cannizzaro About fifty years later, Stanislao Cannizzaro used Avogadro's hypothesis to develop a set of atomic weights for the known elements by comparing the masses of equal volumes of gas.

38. Johann Josef Loschmidt ( Austrian high school teacher ) 1865, calculated the size of a molecule of air, and thus developed an estimate for the number of molecules in a given volume of air.

39. Molar Mass – the mass of one mole of a pure substance One mole of any substance will contain one mole of particles. 6.02 x 10 23  Elements = atoms  Compounds = molecules

40. Calculating Molar Mass 1. multiply atomic mass of each element by number of atoms of that element in the formula (shown by the subscript) 2. find the sum of all the atomic masses --this is formula mass (unit is a.m.u.) 3. express formula mass in grams (unit is g/mol). This is the Molar Mass.