Atomic Structure Review Key

ATOMIC STRUCTURE REVIEW I. Atomic structure: 1. Label each subatomic particle and gives its charge and relative mass: 2. Describe two ways the diagram above is an inaccurate representation of an atom. 3. Determine and number of protons, electrons and neutrons in the following entities: a) Au b) Ca+2 c) P-3 d) 204 Hg & 196 4. Magnesium has isotopes with atomic mass of 24, 25 and 26. Which would be most abundant in a natural sample of this metal? Hg 5. Rubidium has isotopes with atomic masses of 85 (72.17% abundance) and 87 (27.83% abundance). Calculate the average atomic mass of Rb. 6. Why are the very heavy elements radioactive (unstable)? 7. Use your periodic table to write the electron configurations and identify the valence electrons of: a) S c) Ag b) Cs d) Bi 8. Draw the orbital diagrams (all of them!): a) N b) Fe 9. Identify the following on the table below: i) Chemically unreactive ii) A semi-conductor iii) Valence electrons in the 7th energy level. iv) Forms anions of –l charge. v) An alkali-earth metal vi) A radioactive synthetic element. vii) A halogen viii) A transition metals. ix) Most reactive alkali metal x) A noble gas with relatively low ionization energy. 10. Explain the following: a) Fluorine is a smaller atom than oxygen. F has an additional proton – electrons in the 2nd energy level are pulled closer to the nucleus b) Cesium has a lower ionization energy than potassium. The valence electrons of Cs are in the 6th energy level – much further from the nucleus than the valence electron of K (4th energy level). This makes the nuclear attraction to the valence electron of Cs much weaker and therefore easier to remove. c) Oxygen is more electronegative than nitrogen. Oxygen and N both have valence electrons in the 2nd energy level Oxygen is a smaller atom than N because it has more protons which pull on the electrons. This makes oxygen’s nucleus closer to the surface of the atom and more able to attract electrons of other atoms. d) Metallic properties decrease moving across a period. Metallic properties occur when atoms have loosely-held valence electrons. As you go across a period atoms become smaller and hold their valence electrons more tightly and behave less and less like metals. Atomic structure 1 2 I O 3 A T O M I C N U M B E N R B 4 T R A N S I T I O N I 5 Z q U A N T U 6 R A R E E A R T H A T 7 L 8 V a T O M I C R A D I U s 9 a O N 10 E L E C T R O 13 11 E E C A S G 12 A T i V I T E T S R N O O O E M T N R I O G G C P Y M E y 14 T I O N A N I O c 15 P E R I O D 16 A 17 p 18S f S U B S H 19 q U a R ks S O M T I O 20 g L u Y Across Down O N n E L L M 3. The number of protons in an atom. 4. Metals with electrons in the d-subshells 1. The amount of energy required to remove an electron from an atom. 5. A model of electron arrangments that proves electrons can only exist at specific distance from the nucleus. 2. A region where a maximum of 2 electrons can be present. 6. Metals with electrons in the f-subshells. 7. The most distant electrons from the nucleus. 8. Electronegativity goes up as this goes down. 9. This force holds together the nucleus. 10. The ability of an atom to attract electrons in a bond. 11. The number of protons and neutrons in an atom. 13. An atom that has lost electrons. 12. Atoms of the same element with different numbers of neutrons. 14. An atom that has gained electrons. 15. A horizontal row on the periodic table. 18. A region in an energy level where electrons are likely to be. 19. Particles that make-up protons and neutrons. 20. Carrier of the nuclear strong force. 16. Carrier of the electromagnetic force. 17. A vertical column on the periodic table.