CHAPTER 3. Atoms: The Building Blocks of Matter

Transcription

1 CHAPTER 3 Atoms: The Building Blocks of Matter

2 Origins of the Atom Democritus: Greek philosopher (460 BC BC) Coined the term atom from the Greek word atomos Democritus believes that atoms were indivisible and indestructible.

3 Let s Get Ready to Rumble The idea of the atom was met with great skepticism, especially among great thinkers. The most vocal critic of Democritus s idea was Aristotle. Aristotle s chief argument was that there was no proof of the existence of atoms. Democritus s claim was based purely on philosophical argument.

4 Aristotle s Theory of the Atom Aristotle s theory centers around the idea that everything is made up of only 4 elements: earth, wind, fire, and water. THERE ARE NO INDIVISIBLE PARTICLES!

5 Bringing Atoms Back! During the 17th century, scientists quietly revive the idea of the atom. One of the chief supporters include Sir Isaac Newton. At the time, most scientists are concerned with trying to explain the properties of gases.

6 Detour: The Three Basic Laws of Chemistry In their quest to discover proof of atoms, scientists began to propose three basic laws that explain all of the behavior in chemistry. Law of Conservation of Mass (or Matter) Law of Definite Proportions (or Constant Composition) Law of Multiple Proportions

7 Law of Conservation of Mass (or Matter) 1789: French chemist Antoine Lavoisier discovers that during an experiment involving red mercury oxide that the mass of the oxide before heating was equal to the mass of the newly formed mercury metal and oxygen gas. Matter cannot be created or destroyed.

8 Law of Definite Proportions (or Constant Composition) : French chemist Joseph Proust proposes the Law of Definite Proportions based on results from experiments using copper carbonate. Proust finds that all samples of copper carbonate had the same fixed composition. A chemical compound contains the same elements in the same proportions regardless of sample source or size.

9 Law of Multiple Proportions 1803: John Dalton creates an explanation for the Law of Conservation of Mass and the Law of Definite Proportions. As a result, Dalton creates the Law of Multiple Proportions. If 2 or more different compounds are made up of the same elements, then the ratio of the masses of elements is always a small, whole number.

10 Johnny D and the AT 1808: English schoolteacher John Dalton proposes his explanation of 2 of the 3 basic laws of chemistry. In his explanation, Dalton proposes proof of the atoms existence. Dalton s Atomic Theory has 5 main points. You will have a quiz over Dalton s Atomic Theory on Thursday!

11 Dalton s Atomic Theory 1. All matter is composed of atoms. 2. Atoms of a given element are identical in size, mass, and other properties; atoms of different elements differ in size, mass, and other properties. 3. Atoms cannot be subdivided, created, or destroyed.

12 Dalton s Atomic Theory 4. Atoms of different elements combine in simple, whole-number ratios to form chemical compounds. 5. In chemical reactions, atoms are combined, separated, or rearranged.

13 Does Dalton s Theory Still Hold? Not all portions of Dalton s Atomic Theory are still vallid. Now we know that atoms can be subdivided into smaller subatomic particles such as electrons, protons, and neutrons. We also know that a given element can have atoms with different masses.

14 ATOMIC STRUCTURE

15 Discovery of the Electron 1897: J.J. Thomson uses a cathode ray tube to deduce the presence of a negatively charged particle: the electron.

16 How Heavy Is an Electron? 1916: American scientist Robert Millikan determines the mass of an electron to be 1/1837 the mass of a hydrogen atom. In addition, Millikan discovers that an electron has a negative one unit charge.

17 Is There Anything Else Many conclusions were made after the discovery of the electron. 1. All elements must contain identically charged electrons. 2. There has to be positively charged particles if the atom is neutral. 3. There have to be heavier particles since electrons have very little mass.

18 On Another Note 1886: German physicist Eugen Goldstein discovers the proton, a positively charged particle, using anode rays. 1932: English scientist James Chadwick discovers the neutron, a particle with no charge but with a mass slightly larger than a proton.

19 Thomson s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged pudding, thus calling his model the Plum Pudding model. This model has also been called the Blueberry Muffin Model, Chocolate Chip Cookie Model, or the Pepperoni Pizza Model.

20 Ernest Rutherford s Gold Foil Experiment 1911: Ernest Rutherford, Hans Geiger, and Ernest Marsden fire alpha particles at a thin piece of gold foil.

21 Rutherford s Results Most of the alpha particles passed right through; whereas, a few were either deflected or greatly deflected. Rutherford concluded that the nucleus is small, dense, and positively charged.

22 The Rutherford Atomic Model Based on his experiment, we now know The atom is mostly empty space. All of the atom s mass and positive charge is in the nucleus. The nucleus is composed of protons and neutrons. The majority of the atom s volume is the electron cloud.

23 Counting Subatomic Particles Now that scientists have discovered that atoms can be subdivided into subatomic particles, there was a new problem. How do we count subatomic particles? We use terms like atomic number and mass number to do so.

24 Atomic Number Atoms are composed of identical protons, neutrons, and electrons. How are atoms of one element different from those of another element? Each element contains a particular number of protons. The atomic number of an element is the number of protons in the nucleus. # protons in an atom = # electrons (if the atom is neutral!)

25 Mass Number The mass number of an element is the average atomic mass of an element rounded to a whole number. This number is equal to the number of protons and neutrons in the nucleus. Mass number = 195 Mass number = p + + n 0

26 Element # p # e # n Mass no. Carbon 6 Nitrogen 7 14 Sodium Uranium Radon 136

27 Nuclear Symbols Contain the symbol of the element, the mass number, and the atomic number.

28 Nuclear Symbols Find each of these a) number of protons b) number of neutrons c) number of electrons d) atomic number e) mass number

29 Nuclear Symbols If an element has an atomic number of 34 and a mass number of 78, what is the a) number of protons b) number of neutrons c) number of electrons d) complete nuclear symbol

30 Nuclear Symbols If an element has 91 protons and 140 neutrons, what is the a) atomic number b) mass number c) number of electrons d) complete nuclear symbol

31 Isotopes Dalton was wrong about all atoms of elements of the same type being identical. Atoms of the same element can have different numbers of neutrons. Thus, different mass numbers! These atoms are called isotopes. 1912: English radiochemist Frederick Soddy proposes the idea of isotopes. Frederick Soddy : Soddy wins Nobel Prize in Chemistry for this work

32 Naming Isotopes When referencing isotopes of an element, it is important to indicate which mass number the particular isotope has. We typically name isotopes using their element name along with their mass number. Ex. carbon-12, carbon-14, hydrogen-1, hydrogen-2

33 Isotope Protons Neutrons Electrons Mass number

34 Average Atomic Mass Elements occur in nature as a mixture of isotopes. The percentage of each isotope in the naturally occurring element on Earth is nearly always the same, no matter where the element is found. This percentage is taken into account when determining the average atomic mass of a particular element. The average atomic mass is the weighted average of the atomic masses of all naturally occurring isotopes of an element.

35 Calculating Average Atomic Mass To calculate the average atomic mass of an element, two pieces of information will be needed: percent abundance of each isotope and the atomic mass of each isotope. AAM = (mass1 x %A1) + (mass2 x %A2) + *Percent abundance (%A) must be in decimal form in order to conduct calculation.

36 Calculating AAM Ex. 1 Oxygen has three naturally occurring isotopes, oxygen-16, oxygen-17, and oxygen-18. Oxygen-16 has a percent abundance of %, oxygen-17 has an abundance of 0.038%, and oxygen-18 has an abundance of 0.200%. The atomic masses of the three isotopes are amu, amu, and amu, respectively. Calculate the average atomic mass of oxygen.

37 Calculating AAM Ex. 2 There are three naturally occurring isotopes of neon. Their percent abundances and atomic masses are: neon-20, 90.51%, amu; neon-21, 0.27%, amu; neon-22, 9.22%, amu. Calculate the weighted average atomic mass of neon.

38 Calculating AAM Ex.3 Naturally occurring strontium consists of the following isotopes. Isotope Atomic mass, amu Percent abundance Strontium Strontium Strontium Strontium Calculate the weighted average atomic mass of strontium.

39 Calculating AAM Ex.4 The two naturally occurring isotopes of nitrogen are nitrogen-14, with an atomic mass of amu, and nitrogen-15, with an atomic mass of amu. What are the percent abundances of these isotopes?

40 Relating Mass to Number of Atoms How can we determine the number of atoms in a particular number of grams of a substance? We can use three concepts: the mole, Avogadro s number, and molar mass. These three concepts allow us to relate atoms and mass.

41 The Mole Remember, that the mole is the SI base unit for measuring amount of substance. In particular, one mole is equal to the number of particles as there are atoms in exactly 12 grams of carbon-12. The mole is simply a counting unit like a dozen.

42 Avogadro s Number (il numero d Avogadro) Italian chemist Count Amedeo Avogadro devised a way to count the number of representative particles of a substance. Avogadro s number is the number of representative particles in exactly 1 mole of a pure substance. Avogadro s number = x particles

43 Molar Mass We can also define a mole in terms of the amount of substance that contains Avogadro s number of particles. The mass of one mole of a pure substance is called the molar mass of that substance. Molar mass is measured in grams/mole. The molar mass of an element is numerically equal to the atomic mass of the element.

44 Molar Conversions We can convert between particles, moles, and grams. x NA x molar mass atoms moles grams NA molar mass