Ch. 3: Atoms: The Building Blocks of Matter

1. Ch. 3: Atoms: The Building Blocks of Matter 3.1 The Atom: From Philosophical Idea to Scientific Theory 3.2 The structure of the atom 3.3 Counting Atoms

2. 4.1 An atom is the smallest particle of an element that retains its identity in a chemical reaction. Philosophers and scientists have proposed many ideas on the structure of atoms. The lab technician shown here is using a magnifying lens to examine a bacterial culture in a petri dish. When scientists cannot see the details of what they study, they try to obtain experimental data that help fill in the picture.

3. 3 laws • Law of conservation of mass: Mass can neither be created nor destroyed during ordinary chemical and physical reactions. • Problem 1: When ammonium nitrate (NH4NO3) explodes, the products are nitrogen, oxygen, and water. When 40. g of ammonium nitrate explode, 14 g of nitrogen and 8 grams of oxygen form. How many grams of water form? • Law of definite proportions: A pure compound, whatever its source, always contains definite or constant proportions of the elements by mass. • Ex. 1.000 g of NaCl always contains 0.3934 g of sodium and 0.6066 g of chlorine chemically combined. Sodium chloride has definite proportions of sodium and chlorine, irrespective of the source. • Problem 1: If you have 5.00 g of NaCl, how much of it is sodium? • Problem 2: What percent of NaCl is sodium?

4. 3 laws • Law of multiple proportions: It states when two elements form more than one compound, the masses of one element in these compounds for a fixed mass of the other element are in ratios of small whole numbers. • Ex. Carbon and oxygen form two compounds: carbon monoxide (CO) and carbon dioxide (CO2). Carbon monoxide contains 1.3321 g of oxygen for each 1.000 g of carbon, whereas carbon dioxide contains 2.6642 g of oxygen for each 1.000 g of carbon. In other words, the ratio of oxygen that combines with 1.000 g of carbon is 1:2 (small whole number) • Problem 4: Hydrogen and oxygen can combine to form water (H2O) and hydrogen peroxide (H2O2). The mass of oxygen in 1.00 mole of water is 16.0 g and the mass of oxygen in 1.00 mol of hydrogen peroxide is 32.0 g. What is the ratio of oxygen in the two compounds?

5. 4.1 Early Models of the Atom “Everything is made up of a few simple parts called atomos.” Atomos means “uncuttable” in Greek.He envisioned atomos as small, solid particles of many different sizes and shapes.Democritus’s ideas were limited because they didn’t explain chemical behavior and they lacked experimental support. His ideas were rejected because Aristotle supported *the “earth, air, water, and fire” concept of matter. Democritus’s Atomic Philosophy 400 BC (430?) • Democritus believed that atoms were solid particles that are indivisible and indestructible.

6. John Dalton, 1803 (1808?) • By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory. • Dalton’s atomic theory is described in four parts (following slides): “Each element consists of a particular kind of atom. All atoms of a particular element are identical.” His atoms had a specific mass, size, and chemical behavior.

7. 4.1 Early Models of the Atom • 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element. 1. All elements are composed of tiny indivisible particles called atoms.

8. 4.1 Early Models of the Atom • 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element are never changed into atoms of another element in a chemical reaction. 3. Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds.

9. Summary of John Dalton’s atomic theory as listed by tutorvista website: John Dalton, a British school teacher, published his theory about atoms in the year 1808. His findings were based on experiments and also from laws of chemical combination. Main assumptions or postulates of Dalton • All matter consists of indivisible particles called atoms. • Atoms of the same element are similar in shape and mass, but differ from the atoms of other elements. • Atoms cannot be created or destroyed. • Atoms of different elements may combine with each other in a fixed, simple, whole number ratio to form compound atoms. • Atoms of same element can combine in more than one ratio to form two or more compounds. • Atoms are the smallest unit of matter that can take part in a chemical reaction. http://www.tutorvista.com/content/chemistry/chemistry-i/matter/daltons-atomic-theory.php

10. Continued…Drawbacks of Dalton's atomic theory of matter • The indivisibility of an atom was proved wrong, for, an atom can be further subdivided into protons, neutrons and electrons. However an atom is the smallest particle, which takes part in chemical reactions. • According to Dalton, the atoms of same element are similar in all respects. This is wrong because atoms of some elements vary in their mass and density. Such atoms of the same element having different masses are called isotopes. For example, chlorine has two isotopes having mass numbers 35 a.m.u and 37 a.m.u. • Dalton also said atoms of different elements are different in all respects. This has been proved wrong in certain cases like argon and calcium atoms, which have the same atomic mass of 40. Such atoms of different elements that have the same atomic mass are called isobar. • According to Dalton atoms of different elements combine in simple whole number ratio to form compounds. This is not seen in complex organic compounds like sugar C12H22O11. • The theory completely fails to explain the existence of allotropes. The difference in properties of charcoal, graphite, diamond went unexplained in spite of being made up of same kind of atoms.

11. Continued…tutorvista comments Merits of Dalton's atomic theory • It has enabled us to explain the laws of chemical combination. • Dalton was the first person to recognize a workable distinction between the ultimate particle of an element (atom) and that of a compound (molecule).

12. 4.2 3.2 Structure of the Nuclear Atom Cathode-ray tubes are found in TVs, computer monitors, and many other devices with electronic displays. Three kinds of subatomic particles are electrons, protons, and neutrons.

13. 4.2 Subatomic Particles Electrons In 1897, the English physicist J. J. Thomson (1856–1940) discovered the electron. Electrons are negatively charged subatomic particles. Thomson performed experiments that involved passing electric current through gases at low pressure. The result was a glowing beam, or cathode ray, that traveled from the cathode to the anode.

14. 4.2 Subatomic Particles • A cathode ray is deflected by a magnet. Cathode Ray Tube • A cathode ray is deflected by electrically charged plates. • Thomson concluded that a cathode ray is a stream of electrons. Electrons are parts of the atoms of all elements.

15. Charge and mass of the electron • Thomson’s experiment revealed that the electron has a very large charge for its tiny mass. • In 1909, Robert A. Millikan led to the final determination that the mass of the electron is 1/1837 the mass of the simplest type of hydrogen atom, or 9.109 x 10-31 kg. • He also confirmed that the electron carries a negative electric charge. • Based on this information, it was inferred that because atoms are electrically neutral, they must contain a positive charge to balance the negative electrons, and because electrons have such a small mass, atoms must contain other particles to account for most of the mass of an atom.

16. Review: Properties of Electric Charge Electric charge may be only (+) or (-). Opposite charges attract each other Like charges repel.

17. *History of Electric Charge • Discovered by Ben Franklin in the 1750’s in his famous “kite” experiment. Franklin coined the following terms to help describe his findings: • positive (to indicate an object that donates electric charge) • negative (an object that accepts electric charge) • Plus, minus, battery, conductor, and charge • Unfortunately, Franklin decided to describe electricity as moving from (+) to (-). Later, it was discovered (by Thomson) that electricity is the motion of negative charges (electrons), which move from the (-) terminal on a battery to the (+) one. • Today, physicists and physics students still use the convention that electricity flows from (+) to (-), even though they know that electrons flow the opposite way!

18. 4.2 Subatomic Particles In 1886, Eugen Goldstein (1850–1930) observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. He concluded that they were composed of positive particles. Such positively charged subatomic particles are called protons. • In 1932, the English physicist James Chadwick (1891–1974) confirmed the existence of yet another subatomic particle: the neutron. • Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton.

19. 4.2 Subatomic Particles Table 4.1 summarizes the properties of electrons, protons, and neutrons.

20. 4.2 The Atomic Nucleus How can you describe the structure of the nuclear atom? Plum pudding model-- J.J. Thomson and others supposed the atom was filled with positively charged material and the electrons were evenly distributed throughout. Rutherford atomic model--This model of the atom turned out to be short-lived, however, due to the work of Ernest Rutherford (1871–1937).

21. 4.2 The Atomic Nucleus Rutherford’s Gold-Foil Experiment In 1911, Rutherford and his coworkers at the University of Manchester, England, directed a narrow beam of alpha particles at a very thin sheet of gold foil. • Alpha particles scatter from the gold foil.

22. 4.2 The Atomic Nucleus The Rutherford Atomic Model Ernest Rutherford concluded that the atom is mostly empty space. All the positive charge and almost all of the mass are concentrated in a small region called the nucleus. The nucleus is the tiny central core of an atom and is composed of protons and neutrons. In the nuclear atom, the protons and neutrons are located in the nucleus. The electrons are distributed around the nucleus and occupy almost all the volume of the atom.

23. + + Four Forces in the Nucleus Repulsive: 1. Electrostatic: between p+ (like charges repel) Attractive: 2. “Strong force”: between p+ 3. “Weak force”: between p+ and n 4. Gravity: between all particles Strong > electrostatic > weak > gravity Forces in the nucleus:Same charges repel each other but when two protons are extremely close to each other, there are strong attractive forces between them. These forces also exist between p-n and n-n. These short-range forces between p-p, p-n, and n-n are called as nuclear forces and they keep the nuclear particles together.

24. 4.3 3.3 Counting Atoms The atomic number of an element is the number of protons in the nucleus of an atom of that element. Elements are different because they contain different numbers of protons.

25. Practice #1

26. 4.3 Mass Number The total number of protons and neutrons in an atom is called the mass number. The number of neutrons in an atom is the difference between the mass number and atomic number.

27. 4.3 Isotopes Isotopes are atoms that have the same number of protons but different numbers of neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. • Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons.

29. 13C 207Pb 6 82 Isotopic notation Mass number = p + n (in the nucleus) Element symbol • Atomic number = # p • Elements are put in order of atomic number on the periodic table. A E Z Ex: An atom of carbon with 7 neutrons: An atom of lead with 125 neutrons:

30. 4.3 Mass Number Au is the chemical symbol for gold. How many protons, electrons, and neutrons does a gold atom have? The atomic number is 79. Therefore, there are 79 protons and 79 electrons. The mass number is 197, which is the total number of protons and neutrons. Therefore, 197-79= 118 neutrons.

31. Practice #2

32. 4.3 Atomic Mass • Some elements and their isotopes An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom. It is useful to to compare the relative masses of atoms to a standard reference isotope. Carbon-12 is the standard reference isotope. Carbon-12 has a mass of exactly 12 atomic mass units.

33. 4.3 Atomic Mass The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. A weighted average mass reflects both the mass and the relative abundance of the isotopes as they occur in nature.

34. 4.3 Atomic Mass To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. For example, carbon has two stable isotopes: Carbon-12, which has a natural abundance of 98.89% and a mass of 12.000 amu Carbon-13, which has a natural abundance of 1.11% and a mass of 13.003 amu. Silver is found in two isotopes with atomic masses 106.9041 and 108.9047 amu, respectively. The first isotope represents 51.82% and the second 48.18%. Determine the average atomic mass of silver. (106.9041)(.5182)= 55.398 (108.9047)(.4818)= 52.470 55.398 + 52.470 = 107.868 amu

35. Relating Mass to Numbers of Atoms • The mole: One mole of something contains 6.022 x 1023 units of that substance. Ex. One mole of Hydrogen atoms contains exactly 6.022 x 1023 atoms of Hydrogen. • Avogadro’s number: It is defined as the number of particles in one mole of any pure substance. This number is found to be 6.022 x 1023. • Molar mass: Mass of one mole of a substance. Molar mass can be calculated by expressing the mass in amu or g. The molar mass of C = 12 g

36. 10.1 What is a Mole?

39. Converting numbers of atoms to moles On your calculator be sure to put the denominator in parentheses or divide by both 6.02 and 1023

42. Converting moles to number of atoms

43. 10.1 The Mass of a Mole of an Element • The atomic mass of an element expressed in grams is the mass of a mole of the element. • The mass of a mole of an element is its molar mass. • Ex. On the periodic table, carbon has a mass of 12.011 amu. We say 12 grams per mole or just 12 grams.

44. 10.1 The Mass of a Mole of an Element One molar mass of carbon, sulfur, mercury, and iron are shown.

45. 10.1 The Mass of a Mole of an Element

46. 10.1 The Mass of a Mole of a Compound • To calculate the molar mass of a compound, find the number of grams of each element in one mole of the compound. Then add the masses of the elements in the compound. Substitute the unit grams for atomic mass units. Thus 1 mol of SO3 has a mass of 80.1 g.

47. 10.1 The Mass of a Mole of a Compound Molar Masses of Glucose, Water, and Paradichlorobenzene

50. Determining molar mass phosphorus molar mass = 30.974 gchlorine molar mass = 35.453 g 1 mole of phosphorus = 30.974 g3 mole of chlorine = 3 x 35.453 = 106.359 gmolar mass of PCl3 = 30.974 + 106.359 = 137 g/mol