2. Section 4-1 Section 4.1 Early Ideas About Matter (cont.) Dalton's atomic theory The ancient Greeks tried to explain matter, but the scientific study of the atom began with John Dalton in the early 1800's.

3. Section 4-1 The History of Atomic Theory Many ancient scholars believed matter was composed of such things as earth, water, air, and fire. Many believed matter could be endlessly divided into smaller and smaller pieces.

4. Section 4-1 The History of Atomic Theory Democritus (460–370 B.C. ) was the first person to propose the idea that matter was not infinitely divisible, but made up of individual particles called atomos. Aristotle (484–322 B.C. ) disagreed with Democritus because he did not believe empty space could exist. Aristotle’s views went unchallenged for 2,000 years until science developed methods to test the validity of his ideas.

5. Section 4-1 The History of Atomic Theory

6. Late 18th Century Lavoisier proposes the Law of conservation of mass and Proust proposes the Law of constant composition. Early 19th Century Using the previously unconnected ideas above, John Dalton formulates his Atomic Theory. The History of Atomic Theory

7. Section 4-1 The History of Atomic Theory John Dalton revived the idea of the atom in the early 1800s based on numerous chemical reactions. Dalton’s atomic theory easily explained conservation of mass in a reaction as the result of the combination, separation, or rearrangement of atoms.Dalton’s atomic theory

8. Dalton -1803 Billard Ball Model of Atom Atoms are small, round, solid spheres Atomic Theory (5 parts) Elements are made from tiny particles called atoms. All atoms of a given element are identical. (N.B. see isotopes below). The atoms of a given element are different to those of any other element.

9. Section 4-1 Dalton Atoms of different elements combine to form compounds. A given compound always has the same relative numbers and types of atoms. (Law of constant composition). Atoms cannot be created or destroyed in a chemical reaction they are simply rearranged to form new compounds. (Law of conservation of mass).

10. Section 4-2 The Atom The smallest particle of an element that retains the properties of the element is called an atom.atom To give you an idea of how small an atom is we can look at a model. If we use an orange to represent the size of an atom, on the same scale, the orange would be the size of the Earth. An instrument called the scanning tunneling microscope (STM) allows individual atoms to be seen

11. Section 4-2 The Electron When an electric charge is applied, a ray of radiation travels from the cathode to the anode, called a cathode ray.cathode ray Cathode rays are a stream of particles carrying a negative charge. The particles carrying a negative charge are known as electrons.electrons

12. Section 4-2 The Electron (cont.) This figure shows a typical cathode ray tube.

13. Section 4-2 The Electron (cont.) J.J. Thomson measured the effects of both magnetic and electric fields on the cathode ray to determine the charge-to-mass ratio of a charged particle, then compared it to known values. The mass of the charged particle was much less than a hydrogen atom, then the lightest known atom. Thomson received the Nobel Prize in 1906 for identifying the first subatomic particle—the electron. This proved that Dalton was incorrect…atoms were not the smallest particles.

14. Section 4-2 The Electron (cont.) In the early 1910s, Robert Millikan used the oil- drop apparatus shown below to determine the charge of an electron.

15. Section 4-2 The Electron (cont.) Charges change in discrete amounts— 1.602  10 –19 coulombs, the charge of one electron (now equated to a single unit, 1– ). With the electron’s charge and charge-to- mass ratio known, Millikan calculated the mass of a single electron. the mass of a hydrogen atom

16. Section 4-2 The Electron (cont.) Matter is neutral, it has no electric charge. J.J. Thomson's plum pudding model of the atom states that the atom is a uniform, positively charged sphere containing electrons.

17. Section 4-2 The Nucleus In 1911, Ernest Rutherford studied how positively charged alpha particles interacted with solid matter. By aiming the particles at a thin sheet of gold foil, Rutherford expected the paths of the alpha particles to be only slightly altered by a collision with an electron.

18. Section 4-2 The Nucleus (cont.) Although most of the alpha particles went through the gold foil, a few of them bounced back, some at large angles.

19. Section 4-2 The Nucleus (cont.) Rutherford concluded that atoms are mostly empty space. Almost all of the atom's positive charge and almost all of its mass is contained in a dense region in the center of the atom called the nucleus. nucleus Electrons are held within the atom by their attraction to the positively charged nucleus.

20. Section 4-2 The Nucleus (cont.) The repulsive force between the positively charged nucleus and positive alpha particles caused the deflections.

21. Section 4-2 The Nucleus (cont.) Rutherford refined the model to include positively charged particles in the nucleus called protons.protons James Chadwick received the Nobel Prize in 1935 for discovering the existence of neutrons, neutral particles in the nucleus which accounts for the remainder of an atom’s mass. neutrons

22. Section 4-2 The Nucleus (cont.) All atoms are made of three fundamental subatomic particles: the electron, the proton, and the neutron. Atoms are spherically shaped. Atoms are mostly empty space, and electrons travel around the nucleus held by an attraction to the positively charged nucleus.

23. ScientistExperimentKnowledge Gained Relating to J. J. Thomson Cathode RayMass/charge ratio of the electron determined Electron MillikanOil Drop Experiment Charge on the electron Electron Rutherford Gold Foil Experiment Nucleus present in atom The nucleus of an atom and the proton Chadwick NA Neutrons present in nucleus Neutrons Atomic Structure

24. ParticleChargeMass in atomic mass units Position in atom Proton+11Nucleus Neutron01Nucleus Electron1/1836Outside of Nucleus

25. Section 4-3 Section 4.3 How Atoms Differ Explain the role of atomic number in determining the identity of an atom. Define an isotope. Explain why atomic masses are not whole numbers. Calculate the number of electrons, protons, and neutrons in an atom given its mass number and atomic number.

26. Section 4-3 Section 4.3 How Atoms Differ (cont.) atomic number isotopes mass number The number of protons and the mass number define the type of atom. periodic table: a chart that organizes all known elements into a grid of horizontal rows (periods) and vertical columns (groups or families) arranged by increasing atomic number atomic mass unit (amu) atomic mass

27. Section 4.3 Subatomic Particles Most of the atom’s mass. NUCLEUS ELECTRONS PROTONS NEUTRONS NEGATIVE CHARGE POSITIVE CHARGE NEUTRAL CHARGE ATOM QUARKS Atomic Number equals the # of... equal in a neutral atom

28. Section 4-3 Atomic Number Each element contains a unique positive charge in their nucleus. The number of protons in the nucleus of an atom identifies the element and is known as the element’s atomic number.atomic number

29. Section 4-3 Isotopes and Mass Number All atoms of a particular element have the same number of protons and electrons but the number of neutrons in the nucleus can differ. Atoms with the same number of protons but different numbers of neutrons are called isotopes. isotopes

30. Section 4-3 Isotopes and Mass Number (cont.) The relative abundance of each isotope is usually constant. Isotopes containing more neutrons have a greater mass. Isotopes have the same chemical behavior. The mass number is the sum of the protons and neutrons in the nucleus.mass number

31. Atomic & Mass Numbers Mass Number = protons + neutrons Atomic Number = protons only # of Neutrons = mass # - atomic # Mass Number –always a whole number –NOT on the Periodic Table! © Addison-Wesley Publishing Company, Inc.

32. Section 4.3 Isotopes Mass # Atomic # Nuclear symbol: Hyphen notation: carbon-12

33. 4.3 Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen -8 -3342 - 3115 88 16 Arsenic753375 Phosphorus1531 16

34. Section 4.3 Isotopes Chlorine-37 –atomic #: –mass #: –# of protons: –# of electrons: –# of neutrons: 17 37 17 20

35. Section 4-3 Mass of Atoms One atomic mass unit (amu) is defined as 1/12 th the mass of a carbon-12 atom.atomic mass unit One amu is nearly, but not exactly, equal to one proton and one neutron. 12 C atom = 1.992 × 10 -23 g

36. Section 4-3 Mass of Atoms (cont.) The atomic mass of an element is the weighted average mass of the isotopes of that element.atomic mass

37. Average Atomic Mass All periodic tables have atomic mass numbers that are not integers. For example, the atomic mass of Cl is often quoted on periodic tables as 35.5, and may be represented thus, 35.5 Cl 17. This does not mean that there are 17 protons, 17 electrons and 18.5 neutrons in an atom of chlorine. It is not possible to have a fraction of a neutron in an atom.

38. Average Atomic Mass The non-integer values mean that there is more than one isotope of chlorine that exists in nature, in the case of chlorine, 35 Cl and 37 Cl. A quick calculation shows that these two species have the same number of protons and electrons, but different, whole numbers of neutrons (18 and 20 respectively). That is, they are isotopes of one another. These isotopes happen to exist naturally in the approx. abundance, 35 Cl, 75 % and 37 Cl, 25 %.

39. Average Atomic Mass

40. Avg. Atomic Mass EX 1: Calculate the avg. atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. 16.00 amu

41. Avg. Atomic Mass EX 2: Find chlorine’s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. 35.40 amu

42. Example 3: A naturally occurring sample of an element consists of two isotopes, one of mass 85 and one of mass 87. The abundance of these isotopes is 71% and 29%. Calculate the average atomic mass of an atom of this element.

43. Molecules Molecules are formed when a definite number of atoms are joined together by chemical bonds. A molecule can consist of the atoms of only one element, or the atoms of many different elements but always in a fixed proportion.

44. Molecules Molecules are usually formed between non-metal elements. Formula show the number of each type of atom present written as subscripts. The lack of a subscript means only one of that type of atom is present.

45. Examples H 2 O = Water = Compound NH 3 = Ammonia = Compound N 2 = Nitrogen = Element/Compound H 2 = Hydrogen = Element/Compound

46. Ions Atoms have equal numbers of protons and electrons and consequently have no overall charge. When atoms lose or gain electrons, the proton:electron numbers are unbalanced causing the particles to become charged. These charged particles are called ions.

47. Ions Positive ions (where the number of protons is greater than the number of electrons) are called cations, Negative ions (where the number of electrons is greater than the number of protons) are called anions. An ion made up of only one type of atom is called a monatomic ion; An ion made up from more than one type of atom is called a polyatomic ion.

48. Examples Substance Formula Cation or Anion Description Sodium Ion Na + CationMonatomic Chloride Ion Cl - AnionMonatomic Carbonate Ion CO 3 2- AnionPolyatomic Ammonium Ion NH 4 + CationPolyatomic

49. Ions Metals tend to form cations. The metals in groups 1, 2 and 13 tend to lose 1, 2 and 3 electrons respectively to form positive ions with 1+, 2+ and 3+ charges respectively. The transition metals form cations with various positive charges and are much less easy to predict.

50. Non-metals tend to form anions. The non- metals in groups 17, 16 and 15 tend to gain 1, 2 and 3 electrons respectively to form negative ions with 1-, 2- and 3- charge respectively.

51. Examples Consider sodium (i) What is the charge on an ion of sodium? (ii) What is the charge on a sodium nucleus? (iii) What is the charge on a sodium atom? (iv) What is the atomic number of sodium? (v) How many protons in the sodium nucleus? (vi) How many protons in the sodium atom?

52. For the potassium ion, 39 K with a charge of +1; (i) How many electrons, protons and neutrons are present? (ii) Which of these particles has the smallest mass? Examples

53. An ion has a net charge of -1. It has 18 electrons and 20 neutrons. (i) What is its isotopic symbol? (ii) What is its atomic number? (iii) What is its mass number? (iv) What is the charge on the nucleus? (v) How many protons are present?