Chemistry Chapter 4

1. Chemistry Chapter 4 The Structure of the Atom

2. 4.1 Early Ideas About Matter • Objectives: • 1. Compare and contrast the models of Democritus, Aristotle and Dalton • 2. Understand how Dalton’s theory explains conservation of mass

3. Democritus • Definition: An atom is the smallest particle of an element that retains its identity in a chemical reaction • Democritus believed that atoms were indivisible and indestructible • His approach was not based on scientific method

4. Aristotle • Aristotle did not believe empty space could exist • He discredited the ideas of Democritus because Democritus could not explain what held atoms together • Aristotle said that all matter was composed of air, water, earth & fire • These ideas went unchallenged for 2000 years

5. Dalton • Dalton used scientific method and transformed Democritus’s idea into scientific theory • There are 6 points to his theory (check 4.1 in your book) • Here is some of Dalton’s theory: • 1. all elements (matter) are composed of atoms • 2. atoms of the same element are identical (been disproven)

6. 3. Atoms are indivisible & indestructible (been disproven) • 4.Atoms of any element are unique from atoms of any other element • 5.Elements can combine in whole number ratios to form compounds • 6. Chemical reactions occur when atoms are separated, joined, or rearranged • Also, Atoms of one element are never changed into atoms of another element in a chemical reaction

7. Dalton and Conservation of Mass • Recall that mass is always conserved in chemical reactions • Using Dalton’s theory the number of atoms of each type is the same before and after reacting • Dalton’s experiments provided evidence and explanation of the composition of chemical compounds • Conservation of mass led to acceptance of Dalton’s theory

8. Instruments for observing atoms • An atom is the smallest particle of an element that still retains all the properties of that element • Atoms are very small: a copper penny would have roughly 2.4 x 1022 atoms (unfathomably small !!!) • The radius of atom ranges from 5 x 10-10 m to 2 x 10-10 m • Despite their small size atoms are observable with instruments such as a scanning tunneling microscope

9. 4.2 Defining the Atom • Objectives: • 1. Define atom • 2. Distinguish between subatomic particles in terms of relative charge and mass • 3. Describe the structure of the atom, including the locations of the subatomic particles

10. 3 Types of Subatomic Particles • Definition: an atom is the smallest particle of an element that retains the properties of the element • Atoms formed from 3 types of subatomic particles: electrons, protons, and neutrons • Electrons have a negative charge, weigh 1/1840th of protons and neutrons • Symbol is e-

11. Protons have a positive charge and weigh 1840 times as much as electrons • Proton symbol is p+ • Neutrons have no charge, weigh nearly the same as protons • Neutron symbol is n0

12. Thomson, Rutherford & atomic structure • J.J. Thomson- did cathode ray experiments that showed charged particles exist in atoms • Thomson invented the plum pudding model • The model consisted of a spherically shaped atom composed of a uniformly distributed positive charge • The negatively charged electrons existed in the middle of the of the positive charges

13. Rutherford & Thomson were contemporaries • Rutherford set out to disprove Thomson’s model • Rutherford-alpha particle experiments: • 1. Atoms contain a nucleus with protons and neutrons, electrons orbit around nucleus • Definition: nucleus- tiny, central core of atom composed of p+ and n0

14. 2. Atoms composed mainly of empty space • In the nuclear atom, the protons and neutrons are located in the nucleus • The electrons are distributed around the nucleus and the nucleus occupies almost all the volume of an atom

15. 4.3 How atoms differ • Objectives: • 1. Explain the role of atomic number in determining identity of an atom • 2. Calculate the number of electrons, protons & neutrons in an atom, given its mass number and atomic number • 3. Define isotope • 4. Explain why atomic masses are not whole numbers

16. Atomic Number • Elements differ because they have a different number of protons in the nucleus • Protons determine atom identity • How many protons is referred to as atomic number • Given as the top number in the periodic table • In an atom, the number of electrons is equal to the number of protons

17. Definition: Mass Number is the total weight of an atom • Mass number equals the number of protons plus the number of neutrons mass # = #p+ + # n0 • Weight of electrons ignored (1/1840th) • Given as the bottom number in the periodic table • Number of neutrons is the difference between mass number and atomic number • # n0 = mass # - atomic number (#p+)

19. An atom’s composition can be found if the atomic number or mass number is known

20. Why Use the Periodic Table • Definition: a periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties • A periodic table allows you to easily compare the properties of one element (or group of elements) to another element (or group of elements)

21. Definition: a period is each horizontal row of the periodic table • There are seven rows or periods in the modern periodic table • Definition: a group or family is each vertical column of the periodic table • Each group is identified by a number and letter A or B • There are a total of 18 groups, 8 As and 10 Bs

22. Isotopes • Definition: Atoms of the same element with a different number of neutrons are isotopes • Because isotopes of an element have different numbers of neutrons, they also have different mass numbers • Mass number is different, same element • Written with number of neutrons behind, carbon-12, carbon-13 • Or with superscripts for neutrons (13C) and subscripts for protons (6C) put together they are written as (613 C)

23. Actual masses of protons, neutrons and electrons is really small • More useful to talk about relative masses • A reference was needed • Carbon-12 chosen as standard • Carbon-12 was assigned 12 Atomic Mass Units • Definition: atomic mass unit (amu) is one twelfth the mass of a carbon-12 atom • One mass unit is given for each proton and neutron

24. Calculating Atomic Mass: Why Atomic Masses are not Whole Numbers • Most elements in nature are found as mixtures of isotopes • Definition: relative abundance is how often you encounter isotopes in nature and is expressed as a percent • Relative abundance of isotopes determines an averaged mass number • Definition: atomic mass of an element is the weighted average mass of the atoms in a naturally occurring sample of the element • Atomic mass accounts for the isotopes of an element as it occurs in nature

26. To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, then add the products together • Problem: Calculate the atomic mass of bromine. The 2 isotopes of bromine have atomic masses and relative abundances of 78.92 amu (50.69%) and 80.92 amu (49.31%) • A: 79.91 amu

27. 4.4 Unstable Nuclei and Radioactive Decay • Objectives: • 1. Explain the relationship between unstable nuclei and radioactive decay • 2. Characterize alpha, beta and gamma radiation in terms of mass and charge

28. The Relationship • Definition: radioactivity is spontaneously emitted radiation produced by some substances • Definition: radiation is the rays and particles emitted by those substances • Definition: nuclear radiation is a reaction that involves a change in an UNSTABLE atomic nucleus

29. Most radioactive atoms tend to be heavier elements with a large nucleus making it emit radiation because their nucleus is unstable • Definition: radioactive decay is a spontaneous process in which unstable nuclei lose energy by emitting radiation • 200 years ago, scientists began investigating radiation by directing it between electrically charged plates

30. Definition: alpha radiation is the radiation that was deflected toward the negatively charged plate • It consists of alpha particles • Alpha particles are composed of 2 neutrons and 2 protons, & no electrons • Therefore an alpha particle carries a charge of 2+ & is the nucleus of a helium (He) atom

31. Definition: beta radiation is the radiation that was deflected toward the positively charged plate • Each beta particle is an electron with a 1- charge & is an electron which is emitted • Definition: gamma rays or gamma radiation is high energy radiation with no mass • It is denoted by the Greek letter gamma which looks like this:

32. The main factor in determining nuclear stability is the ratio of neutrons to protons • Too many or too few neutrons are unstable and cause the nucleus to lose energy through radioactive decay until a stable nucleus is formed • Alpha and beta particles are emitted until a stable nonradioactive nucleus is formed

33. Alpha radiation is the least harmful, can burn skin & is stopped by paper or clothing • is not harmful unless ingested or breathed in • The electron of a beta particle is stopped by aluminum & is more penetrating than alpha • Gamma radiation is pure energy which is very penetrating & harmful

34. Alpha, beta & gamma equations • Examples of alpha: • Examples of beta: • Gamma will emit either alpha or beta particle + gamma rays