Chapter 4: Atomic Structure

1. Chapter 4: Atomic Structure

2. 4.1 Defining the Atom • An Atom is the smallest particle of an element that retains its identity in a chemical reaction. • The Greek Philosopher Democritus, was among the First to suggest the existence Of atoms.

3. Democritus believed that atoms were indivisible and indestructible. • Although Democritus’s ideas agreed with later scientific theory, they did not explain chemical behavior. • They also lacked experimental support.

4. Dalton’s Atomic Theory • By experimental methods, Dalton transformed Democritus’s ideas on atoms into scientific theory. • Dalton studied the ratios in which elements combine in chemical reactions. Based on the results of his experiments, Dalton formulated hypotheses and theories to explain his observations.

5. Dalton’s Atomic Theory • All elements are composed of tiny indivisible particles called atoms. • Atoms of the same element are identical. The atoms of any one element are different from those of any other element. • Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. • Chemical reactions occur when atoms are separated, joined, or rearranged.

6. Sizing up the Atom…. • Copper atoms are very small. • A copper coin the size of a penny contains about 2.4 x 10^22 atoms. By comparison, Earth’s population is only about 6 x 10^9 people. • If you could line up100,000,000 Copper atoms side by side, they would produce a line only 1 cm long.

7. 4.2 Structure of the Nuclear Atom • Much of Dalton’s atomic theory is accepted today. One important change however, is that atoms are now known to be divisible. • Three kinds of subatomic particles are electrons, protons, and neutrons.

8. Electrons • In 1897, English physicist J.J. Thomson discovered the Electron. • Electrons are negatively charged subatomic particles. • Thomson called these particles corpuscles; later they were named electrons.

10. Robert A. Millikan (1868-1953) • Carried out experiments to find the quantity of charge carried by an electron. • An electron carries exactly 1 negative charge and its mass is 1/1840 the mass of a hydrogen atom.

11. You need to understand…. • Atoms have NO CHARGE!! • They are electrically neutral.

12. Next…….. • Electric charges are carried by particles of matter.

13. Third…… • There are NO fractions of charges, only whole number multiples!!!

14. And Finally……. • When a negative particle combines with a positive particle an electrically neutral particle is formed. +1 + -1 = 0

15. Eugene Goldstein • Observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. • He called these rays canal rays and concluded that they were composed of positive particles. • These positive particles are called Protons.

16. James Chadwick • In 1932, an English Physicist confirmed the existence of yet another subatomic particle: The Neutron. • A neutron is a subatomic particle with nocharge but with a mass nearly equal to that of a proton.

17. In Thomson’s atomic model, known as the plum pudding model, electrons were stuck into a lump of positive charge, similar to raisins stuck in dough. Plum Pudding Model

18. This model of the atom turned out to be short lived, however, due to the groundbreaking work of Ernest Rutherford (1871-1937), a former student of Thomson.

19. Rutherford’s Gold Foil Experiment

20. Rutherford’s Gold Foil Experiment • Here is a video of what happened!!!

21. Rutherford Atomic Model • He proposed that the atom is mostly empty space, thus explaining the lack of deflection of most of the alpha particles. • He concluded that all the positive charge and almost all the mass are concentrated in a small region. • He called this region the Nucleus. • The nucleus is the tiny central core of an atom and is composed of protons and neutrons.

22. Key Concept • 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. Although it was an improvement over Thomson’s model of the atom, Rutherford’s model turned out to be incomplete. • In Chapter 5, we will learn how it was revised.

24. Elements are different because they contain different numbers of protons. 4.3 Distinguishing Among Atoms

25. Atomic Number • The atomic number of an element is the number of protons in the nucleus of an atom of that element.

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

27. If you know the atomic number and mass number of an atom of any element, you can determine the atom’s composition. How do you calculate mass number?

28. 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 have different mass numbers.

29. Atomic mass unit is defined as one twelfth of the mass of a carbon-12 atom. In nature most elements occur as a mixture of two or more isotopes. Isotopes of an element has a fixed mass and a natural percent abundance. Isotopes Cont…

30. Relative Abundance • 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.

31. Determine Atomic Mass • # stable isotopes • Mass of each isotope • Natural % abundance of each isotope Atomic Mass = Mass of Isotope x natural abundance = add them together Look on page 117 for practice problems!!

32. Each horizontal row of the periodic table is called a PERIOD. The Periodic Table – A Preview

33. Each vertical column of the periodic table is called a group. Elements within a group have similar chemical and physical properties. Periodic Table Cont….

34. Lets Review!!! • Atoms are the smallest particle of an element that retains its identity in a chemical reaction. • Democritus believed that atoms were indivisible and indestructible. • Dalton’s atomic theory states that • All elements are composed of atoms • Atoms of different elements combine to form compounds • Atoms of the same element are identical • Chemical reactions occur when atoms are separated, joined, or rearranged.

35. Review Cont…. • There are 3 subatomic particles • Proton • Identifies the atom • Positive charge • Electron • Determines chemical reactions • Negative Charge • Neutron • Determines Isotope • Neutral Charge

36. Review Cont… • Thomson atomic model, known as the plum pudding model, electrons were stuck into a lump of positive charge. • Rutherford’s Gold Foil Experiment: He proposed that atoms are mostly empty space. • The nucleus is the tiny central core of an atom made of protons and neutrons. Electrons surround the nucleus. • Atomic number is the number of protons. • Mass number is the # of protons and neutrons. • Isotopes are atoms that have the same number of protons but a different number of neutrons.

37. Review Cont… • Atomic mass unit is defined as one twelfth of the mass of a carbon-12 atom. • The atomic mass of an element is the weighted average mass of the atoms in a naturally occurring sample of the element. • To calculate the atomic mass of an element, multiply the mass by its natural abundance and then add the products.

38. Review Cont… • Periodic Table allows you to easily compare the properties of one element to another element. • Groups go down!! • Periods all across!! • Elements within a group has similar physical and chemical properties.