Section 4–2: The Structure of an Atom

1. Section 4–2:The Structure of an Atom Coach Kelsoe Physical Science Pages 108–112

2. Objectives • Identify three subatomic particles and compare their properties. • Distinguish the atomic number of an element from the mass number of an isotope, and use these numbers to describe the structure of atoms.

3. Properties of Subatomic Particles • By 1920, Rutherford had seen evidence for the existence of two subatomic particles and had predicted the existence of a third particle. • There are three subatomic particles – • Protons • Electrons • Neutrons

4. Protons • Rutherford determined that the amount of positive charge varies among elements. • Each nucleus must contain at least one particle with a positive charge. These particles are called protons. • A proton is a positively charged subatomic particle that is found in the nucleus of an atom. Each proton is assigned a charge of 1+.

5. Electrons • The particles that Thomson detected were later named electrons. • An electron is a negatively charged subatomic that is found in the space outside the nucleus. Each electron has a charge of 1-.

6. Neutrons • In 1932, the English physicist James Chadwick designed an experiment to show that neutrons exist. He concluded that the particles he produced were neutral because a charged object did not deflect their paths. • A neutron is a neutral subatomic particle that is found in the nucleus of an atom. Its mass is almost exactly equal to that of a proton.

7. Comparing Subatomic Particles • Protons, electrons, and neutrons can be distinguished by mass, charge, and location in an atom. • Protons and neutrons have almost the exact same mass, but electrons are 1/2000 the size of protons. Although an electron is smaller in mass, they have an equivalent negative charge to a proton’s positive charge.

8. Atomic Number • The atoms of any given element always have the same number of protons. There are always 3 protons in lithium and 4 in beryllium. • The atomic numberof an element equals the number of protons in an atom of that element.

9. Atomic Number • Atoms of different elements have different numbers of protons. This is what gives atoms their identity! • Each positive charge in an atom is balanced by a negative charge because atoms are neutral. So the atomic number of an element also equals the number of electrons in an atom. Lithium has 3 electrons; beryllium has 4.

10. Mass Number • The atomic number tells you the number of protons in the nucleus of an atom, but doesn’t tell you how many neutrons there are. • The mass number of an atom is the sum of the protons and neutrons in the nucleus of that atom. • For example, an atom of carbon that has 6 protons and 7 neutrons would have a mass number of 13.

11. Mass Number • If you know the atomic number and mass number of an atom, you can find the number of neutrons by subtracting. • Example: 23 (MN of Na) – 11 (AN of Na) = 12 Number of neutrons = Mass number – Atomic Number

12. Isotopes • In Dalton’s atomic theory, he states that all atoms of a given element are identical. Every atom of a given element has the same number of protons as electrons. • However, atoms of the same element can have different masses due to the number of neutrons!

13. Isotopes • Isotopes are atoms of the same element that have different numbers of neutrons and different mass numbers. • Isotopes of an element have the same atomic number but different mass numbers because they have different numbers of neutrons.

14. Isotopes • For example, there are three types of oxygen: oxygen-16, oxygen-17, and oxygen-18. The number at the end represents the mass number. • All of these oxygen atoms have 8 protons (its atomic number is 8), but the number of neutrons will change: • Oxygen 16  16-8 = 8 neutrons • Oxygen 17  17-8 = 9 neutrons • Oxygen 18  18-8 = 10 neutrons

15. Vocabulary • Proton • Electron • Neutron • Atomic number • Mass number • Isotopes