Chapter 5: Early Atomic Theory and Structure. 5.1 Early Thoughts. In the year 440 B.C., believed that all matter was made of 4 elements (list them):

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1 Chapter 5: Early Atomic Theory and Structure Name: 5.1 Early Thoughts In the year 440 B.C., believed that all matter was made of 4 elements (list them): Around 370 B.C., proposed that all matter was composed of tiny, indivisible particles (the word meant indivisible). This is where the idea of the atom was born. But, there were no experiments to support his claims Dalton s Model of the Atom More than 2,000 years later in 1808, John Dalton proposed an atomic model based on facts and experimental evidence: 1.) Elements are composed of and particles called atoms. 2.) Atoms of the same element are identical in mass and size. 3.) Atoms of different elements have different and. 4.) 2 or more atoms of different elements can combine to form 5.) Law of Definite Proportion: 6.) Law of Multiple Proportions: Ex: * Modifications of Dalton s Theory: overall model still holds true, but later investigations showed the following: Atoms are composed of subatomic particles (atoms are in fact divisible). Not all atoms of a specific element have the same mass ( isotopes : see Sect 5.8). Under special circumstances, atoms can be decomposed (atoms are not indestructible) Subatomic Parts of the Atom An atom is extremely small! But, there are even smaller parts which make up an atom... called subatomic particles. Particle Symbol Electrical Charge Mass (g) Discovery of Subatomic Particles: 1. Electron. Who was responsible for the discovery of the negatively charged electron? In what year? He did this by constructing a piece of equipment called a (or sometimes called a. 1

2 Thomson Model of the Atom: aka What did it look like? (label all parts)?: 2. Proton. Who was responsible for the discovery of the positively charged proton? In what year? 3. Neutron. Who was responsible for the discovery of the neutron? In what year? Ions: positively or negatively charged atom/group of atoms. Positive Ion (also called a ) is formed when Ex: Negative Ion (also called a ) is formed when Ex: 5.7 The Nuclear Atom In 1911, an experiment by called the showed the existence of the nucleus of an atom. Experiment: He shot a stream of positively charged at a thin piece of gold metal foil. He found that most of the particles But, some of the particles Results: Rutherford asked himself what in an atom could have deflected some of the particles? Rutherford knew that the were far too light to deflect these alpha particles. He also knew that like charges repel, so something small, dense, & positively charged in the atom must be deflecting these positively charged alpha particles he called this the. He also concluded that since most of the alpha particles went straight through, most of an atom was. The work of Rutherford and later Chadwick led to the Nuclear Model of the Atom: 2

3 General Arrangement of Subatomic Particles: Each atom consists of a surrounded by in the. Inside the nucleus are the and. This is where the majority of an atom s mass comes from. In a neutral atom, the positive charge of the nucleus (due to the ) is exactly offset by the negative electrons. Therefore, a neutral atom must contain exactly the same number of and. Draw a sketch of the Nuclear Model of the Atom (label all parts with their charges): Define Atomic Number: The atomic # determines the identity of an atom! Ex: every atom with an atomic # of 1 is a hydrogen atom. Define Mass Number: *Atoms of the same element can have different masses because of Isotopes. Define an Isotope: What does Isotopic Notation look like (label the parts)? Examples: Isotopic Notation Atomic Number Mass Number # of Protons # of Neutrons # of Electrons 3

4 5.9 - Atomic Mass The mass of a single atom is too small to measure so a system of relative atomic mass units (amu) was devised. What is used as the standard for atomic masses? Are all of the isotopes of a given element found equally in nature? We use what s called to express how often we can find that isotope in nature. Then we use this information to calculate an average atomic mass. Define Average Atomic Mass: What is the equation to find the Average Atomic Mass? Ex 1: Calculate the average atomic mass of oxygen if its abundance in nature is 99.76% 16 O, 0.04% 17 O, and 0.20% 18 O. Ex 2: Find chlorine s average atomic mass if approximately 8 of every 10 atoms are chlorine-35 and 2 are chlorine-37. 4

5 : Ions Name: 1.) Define an ion: 2.) List the number of protons, electrons, & neutrons for the following: protons electrons neutrons *A positive ion is called a: Ca 20 Ca +2 3.) List the number of protons, electrons, & neutrons for the following: protons electrons neutrons *A negative ion is called an: F 9 F -1 4.) Does the number of protons for an ion differ from its neutral atom? Explain 5.) Does the number of neutrons for an ion differ from its neutral atom? Explain 6.) Does the number of electrons for an ion differ from its neutral atom? Explain 7.) Does the atomic number for an ion differ from its neutral atom? Explain 8.) Does the mass number for an ion differ from its neutral atom? Explain 9.) Describe the difference between a neutral atom and a cation. 10.) Describe the difference between a neutral atom and an anion. 5

6 5.7: Atomic Number and Mass Number Name: Complete the following chart and answer the questions below. Element Name Atomic Number Number of Protons Number of Neutrons Mass Number Number of Electrons carbon hydrogen hydrogen 2 nitrogen cesium tungsten silver chlorine with -1 charge calcium with +2 charge ) How are the atomic number and the number of protons related to each other? 2.) How do the number of protons, number of neutrons, and the mass number relate to each other? 3.) What is the one thing that determines the identity of an atom (that is, whether it is an oxygen atom or a carbon atom, etc.)? 4.) For a neutral atom, how do the number of protons and electrons relate to each other? 6

7 5.9: Isotopes Name: 1.) Atoms are made up of the subatomic particles: protons, neutrons, and electrons. The nuclei of atoms that make up isotopes of an element differ. There are 3 known isotopes of the element hydrogen. List the number of protons, electrons, and neutrons for each of these isotopes below: Protium 1 1 H Deuterium 12 H Tritium 1 3 H protons electrons neutrons 2.) Explain why the atomic mass of hydrogen is g and not a whole number. 3.) Which of the hydrogen isotopes is most abundant in nature? Explain. 4.) Do the atomic numbers of isotopes differ? Explain. 5.) Do the mass numbers of isotopes differ? Explain. 6.) The element neon consists of three isotopes whose masses are, respectively 20.0g, 21.0g, and 22.0g. The abundance of these isotopes is respectively 90.92%, 0.25%, and 8.83%. Just by looking at the percents, what is the approximate atomic mass of neon? 7.) The element boron consists of two isotopes of masses of 10.0g and 11.0g. The average mass of boron is g. Which isotope is more abundant? 7

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