Atoms: the building blocks of matter

1. Atoms: the building blocks of matter Chapter 3 Chemistry chapter 3

2. The atom • The atom – smallest piece of matter that has the properties of an element. • Made of • Protons • Neutrons • Electrons • Each specimen of a specific subatomic particle is the same •  If we split an atom, we no longer have a specific element Chemistry chapter 3

3. Early atomic theory - Democritus • Greek philosopher about 400 B.C. • Gave us the word atom • Atomos - indivisible. • Thought • The world was made of empty space and particles called atoms. • There were different types of atoms for different types of materials. • Theory was not supported by experimental evidence. Chemistry chapter 3

4. Early atomic theory – Aristotle • Aristotle did not believe in atoms • thought matter was continuous • He was very influential, so Democritus’s theory was not accepted for many centuries. Chemistry chapter 3

5. 17th century • People began to express doubts in Aristotle’s theory. • Isaac Newton and Robert Boyle published articles stating their belief in the atomic nature of elements, but they had no proof. • Their theory also had no ability to predict the unknown. Chemistry chapter 3

6. Antoine Lavoisier – late 1700s • Law of conservation of mass • during a chemical change in a closed system, no mass is lost Chemistry chapter 3

7. Joseph Proust – late 1700s to early 1800s • Law of definite proportions • specific substances always contain elements in the same ratio by mass Chemistry chapter 3

8. Law of multiple proportions • Some elements form more than one compound with each other. • If two or more different compounds are composed of the same two elements, then the ratio of their masses always contains small whole numbers Chemistry chapter 3

9. John Dalton – early 1800s • Studied experimental observations of chemical reactions • Proposed explanation of these three laws Chemistry chapter 3

10. Dalton’s Hypothesis • All matter is composed of very small particles called atoms. • All atoms of an element are exactly alike; atoms of different elements are very different. • Atoms cannot be subdivided, created, or destroyed. • Atoms unite with other atoms in simple ratios to form compounds • In chemical reactions, atoms are combined, separated, or rearranged. Chemistry chapter 3

11. Did Dalton’s theory work? • Conservation of mass • the atoms are simply rearranged because they cannot be created or destroyed • Laws of definite and multiple proportions • Only whole atoms can combine, giving small whole numbers in ratios Chemistry chapter 3

12. Gas research • J.L. Gay-Lussac • Under constant temperature and pressure • Volumes of reacting gases and gaseous products are in a ratio of small whole numbers. • Amadeo Avogadro explained Gay-Lussac’s work with Dalton’s theory. • Equal volumes of gases, under the same temperature and pressure, have the same number of molecules. • Helped Dalton’s theory get accepted Chemistry chapter 3

13. Dalton’s theories • Atomic theory and law of multiple proportions have been tested and accepted as correct. • However, there some major exceptions to the rules. • Splitting atoms • Different atoms of the same element Chemistry chapter 3

14. Discussion • Section review on page 69 Chemistry chapter 3

15. Cathode tubes • Anode – positive electrode • Cathode – negative electrode • When the tube is on, cathode rays appear that begin at the cathode and travel to the anode. Chemistry chapter 3

16. Cathode rays and electrons • 1897 – J.J Thomson tested cathode rays and discovered that they were electrons. • Rays turned a paddlewheel – they had mass • Rays deflected by a magnet just like current-carrying wire – they were negatively charged • He determined the ratio of the electron’s charge to its mass. Chemistry chapter 3

17. Charge on an electron • Robert Millikan’s famous oil drop experiment. • Tiny oil drops fell through a chamber • gravitational force offset by applying an opposing electrical force. • Charge on oil drops determined • This charge was always a whole number multiple of one small charge Chemistry chapter 3

18. Charge on an electron • This small charge was the charge on one electron. • This is now the standard unit of negative charge (1-). It can be written e-. • e- can also represent an electron Chemistry chapter 3

19. Mass of an electron • Using Thomson’s ratio and Millikan’s charge, determined to be 9.1 x 10-31 kg • It was found that it’s mass is only 1/1837 the mass of the lightest atom known – the hydrogen atom. • Most of the mass must be somewhere else • Since atoms are neutral, there must be some positive charge Chemistry chapter 3

20. Thomson’s plum pudding model • In this model, the raisins were the electrons and the pudding was the positive charge. • Sort of like chocolate chip cookie dough. • The chips are the electrons and the dough is the positive charge. • Explained the experiments that had been done so far. Chemistry chapter 3

21. Testing the plum pudding model • See page 72 • fired alpha particles at a very thin (a few atoms thick) sheet of gold foil. • They expected the particles to go right through because the spread out positive charge in the “pudding” wouldn’t be strong enough to deflect them. Chemistry chapter 3

22. What happened • Most of the particles did go right through without being deflected at all. • Some were deflected at large angles. • Ernest Rutherford explained it: • the positive charge on the atom was concentrated at a small core – now called the nucleus. Chemistry chapter 3

23. The atom as we now “know” it • The nucleus contains all of the positive charge and most of the mass. • The negatively charged electrons have very small mass and are located around the nucleus in the electron cloud. • Most of an atom is empty space. Chemistry chapter 3

24. Protons • same charge as an electron; opposite sign. • standard unit of positive charge (1+) • Much larger mass than the electron: 1.67 x 10-27 kg • The number of protons determines the atom’s identity. Chemistry chapter 3

25. Neutrons • Weren’t discovered until the 1930s. • Neutral – no charge – harder to detect • Slightly more mass than a proton: 1.68 x 10-27 kg Chemistry chapter 3

26. Nuclear or Strong Force • The force that holds protons and neutrons together. • It is effective only for very short distances – about 10-15 m. Chemistry chapter 3

27. Dalton’s theory • Dalton thought that atoms were indivisible • discovery of electrons, protons, and neutrons did not fit with his theory. • Led to major revisions in atomic theory Chemistry chapter 3

28. Isotopes • Thomson discovered what seemed to be two kinds of neon atoms. • Same chemical properties; different masses. • Atoms of the same element that differ in mass are called isotopes. • Have the same number of electrons and protons but different number of neutrons. Chemistry chapter 3

29. Atomic number • Number of protons in an atom • Atoms are electrically neutral,  the number of electrons must equal the number of protons. • The number of protons determines the identity of the atom and the number of neutrons determines the isotope. Chemistry chapter 3

30. Modification of Dalton’s theory • All atoms of an element contain the same number of protons but can contain different numbers of neutrons. • So we have to use average mass of an atom. Chemistry chapter 3

31. Nucleons • Particles in the nucleus – protons and neutrons Chemistry chapter 3

32. Mass number • Total number of nucleons : protons plus neutrons • Number of neutrons = mass number minus atomic number Chemistry chapter 3

33. Designating Isotopes • Hyphen notation • Uranium-235 • Carbon-14 • Carbon-12 • The number refers to the mass number Chemistry chapter 3

34. Nuclide • General term for any isotope of any element Chemistry chapter 3

35. Atomic mass units • There must be a standard for all units of measurement. • A Carbon-12 atom with 6 protons and 6 neutrons was chosen as the standard Chemistry chapter 3

36. Atomic mass unit • Defined as 1/12 the mass of that carbon atom. Chemistry chapter 3

37. Average atomic masses • Many elements have an average atomic mass close to the number of nucleons in their nuclei – near whole numbers. • Some don’t – look at Chlorine • The periodic table shows average atomic masses. Chemistry chapter 3

38. Weighted averages • We then use a weighted average to find the average mass of an atom of a given element. • This is called the average atomic mass or just atomic mass. Chemistry chapter 3

39. Finding a weighted average • A class of 25 students took a test. 10 of them got 80%. 12 got 90%. 3 got 100%. What was the average score? • Not 90% - probably less than that. Chemistry chapter 3

40. You try • Neon has two isotopes. Neon-20 has a mass of 19.992 amu and neon-22 has a mass of 21.991 amu. In any sample of 100 neon atoms, 90 will be neon-20 and 10 will be neon-22. Calculate the average atomic mass of neon. • 20.192 amu Chemistry chapter 3

41. You try • Compute the average atomic mass of silver, if 51.83% of the silver atoms occurring in nature have mass 106.905 amu and 48.17% of the atoms have mass 108.905 amu. • 107.9 amu Chemistry chapter 3

42. The Mole • SI unit for amount of substance • Abbreviated mol • A counting unit • 6.022 x 1023 particles • Avogadro’s number • Based on carbon-12, 12 g of C-12 contains 1 mol of atoms Chemistry chapter 3

43. Molar mass • The mass of 1 mol of a pure substance • g/mol • Numerically equal to the atomic mass in amu • On the periodic table the number with a decimal is the atomic mass in amu AND the molar mass in g/mol Chemistry chapter 3

44. conversions • Grams to moles or moles to grams • Use the molar mass Chemistry chapter 3

45. Example • What is the mass in grams of 5.60 mol of sulfur? Chemistry chapter 3

46. Example • How many moles of carbon are in a sample with a mass of 567 g? Chemistry chapter 3

47. Example • How many atoms of lithium are in a sample with a mass of 76.2 g? Chemistry chapter 3

48. You try • How many moles of rubidium are in 3.01 x 1023 atoms of rubidium? Chemistry chapter 3

49. You try • How many moles are in 0.255 g of zinc? Chemistry chapter 3

50. You try • What is the mass of 1.20 x 1025 atoms of helium? Chemistry chapter 3