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Chapter 4: Atomic Structure (Chem I/Chem IH)

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1 Chapter 4: Atomic Structure (Chem I/Chem IH)

2 Composition of Atoms (see Table 4.1 on p 106)
Protons: positive charge (+1), located in nucleus, are heavy, “p+” Neutrons: no charge(0), located in nucleus, are heavy, “n0” Electrons: negative charge (-1), located outside nucleus “electron cloud”, very light (1/1840 of a proton or neutron), “e-”


4 Charges in an Atom The + charge on a proton is equal to the - charge on an electron. Atoms are neutral (have no overall charge) Key Concept: Therefore, the # of protons = # electrons in an atom.

5 Atomic number Located on periodic table, usually physically largest number determines the identity of the atom. It tells us the number of protons in the atom. It also tells us the number of electrons (b/c an atom is neutral in charge.) Ex: atomic number of carbon, C = 6 Question: how many p+s? How many e-s? Ex: p 111 Problem 4.1: Nitrogen has atomic number of 7. How many protons and electrons in a neutral atom of nitrogen? Slide 2.2

6 Isotopes The number of neutrons can vary from atom to atom in an element. Atoms of the same element w/different #s of neutrons are called ISOTOPES. In order to know how many neutrons in an atom you must be told. Either you are told # of neutrons directly OR Mass number

7 Isotope Symbols & Names
Isotope Name carbon -14 C Mass number 14 6 Atomic number

8 Mass Number The mass number tells you how much mass the atom has.
Since p+ and n0 are the heavy parts, mass # = # of p+’s + n0’s. How to use this to determine composition of atom Ex: p 112, # 4.1: identify p+, n0, e- A) beryllium-9: B) neon-20 C) sodium-23 (sodium is Na)

9 HW: Pp Problems # 17-20

10 QUESTION: If the mass number of a carbon atom is 14,
How many protons? How many electrons? How many neutrons? LET’S PRACTICE! Whiteboard Marker Paper towel

11 Atomic Mass Units Atoms are weighed in a.m.u.
1 a.m.u. is based on the mass of a Carbon-12 atom. it has 6 p+ and 6 n0, 1 a.m.u = 1/12 the mass of a carbon-12 atom.

12 Atomic Mass (definition) Weighted average of all the isotopes of an element. See p 114 of text. calculating atomic mass often located below element symbol on periodic table. Usually a decimal

13 Relative Abundance & Atomic Mass
Sample problem 4.2 p 117 of text Element X has 2 natural isotopes. The isotope with a mass of amu has a relative abundance of 19.91%. The isotope with a mass of amu has a relative abundance of 80.09%. Calculate the atomic mass of this element.

14 Analysis of Problem Isotope: Relative Abundance 10X % 11X 80.09% Convert % to decimal values = =

15 Solution To find out how much each isotope contributes to the element’s atomic mass, multiply the isotope’s mass by its relative abundance amu x = amu amu x = amu

16 Solution ,cont. Now add these two numbers to find the atomic mass for the element, X. amu. Does this make sense? (Should your answer be closer to 10 or 11?)

17 Early Ideas About Matter
Greek philosopher, Democritus (2500 y.a.) proposed matter was made of 4 elements: earth, air, fire, water. (PLEASE NOTE: None of these are elements!) He coined the word “atom” meaning “cannot be broken.”

18 “Modern” Atomic Theory
Proust (1799) discovered water was always 11% hydrogen, 89% oxygen – law of definite proportions Compounds’ components are always in a specific proportion by mass

19 Discovery of Atomic Structure
1800s- scientists thought atom was a tiny solid ball. THEN…JJ Thomson (1897) discovers the electron (e-)

20 JJ Thomson Cathode Ray Experiment
Vacuum tube (no air inside) w/ electrode on each end, attached to a terminal. He send electricity through the tube and saw A bright ray travelling from the negative end (cathode) to the positive end (anode). “cathode rays” picture of cathode ray tube Cathode ray bends toward a positive end of a magnet. “ “ bends away from a negative end of a magnet.

21 Cathode Ray Experiment
Conclusion: 1. The cathode ray was actually ____ charged particles. 2. The atom could not be ________ as scientists had thought, but must contain charged particles.

22 (Section 2):Electrons in Atoms
Energy of Electrons Why electrons don’t crash into the nucleus: they have enough energy to keep them away. Why e-s (usually) don’t fly off of atoms: they have enough attraction to the nucleus to keep them in “orbit.” (Kind of like planets in orbit around the sun.)

23 Discovery of Protons & Neutrons, cont.
In 1910 Thomson discovered that neon atoms have different masses. In 1932, James Chadwick confirms existence of the neutron Conclusion: there must be another particle that has no charge, called a neutron.

24 Dalton’s Atomic Theory, cont.
John Dalton ( ) 1. All matter is made of atoms 2. Atoms are indestructible and can’t be divided 3. All atoms of one element are exactly alike, but different from atoms of other elements.

25 Hypotheses, Theories, Laws (review)
Hypothesis: testable prediction to explain an observation Theory: well tested explanation that explains many observations. May change over time Law: fact of nature observed so often it is accepted as truth. Doesn’t change.

26 Discovery of Protons & Neutrons
Eugen Goldstein realized there was a second ray in the vacuum tube. It bent toward the – end of a magnet. It bent away from the + end of a magnet. Therefore, this ray was made of _____ charged particles.

27 Discovery of Nucleus 1909-scientists now believe the atom is like chocolate chip cookie dough (see Fig 2.8 p 63 of text) 1911-Rutherford’s Gold Foil Experiment Shot “alpha Particles” (helium nuclei) at gold foil. Hypothesis: they would pass through unaffected. Data: most did pass through Some were deflected Others bounced straight back!

28 The Nuclear Model of the Atom
Conclusion: 1. Atoms are nearly all empty space! 2. Atoms have small, densely packed central nucleus

29 Energy of Electrons (cont.)
(Don’t write this!) DISCUSS WITH YOUR NEIGHBOR: You are an electron. If you have a lot of energy, will you stay close to the nucleus or will you move further from it? Answer: you may still stay in “orbit” but you will be able to move further away from the nucleus.

30 Bohr’s Model of Atom Neils Bohr studied w/Rutherford
His model is also called the planetary model He discovered that e-s could only exist at certain distances from the nucleus.

31 Bohr’s Model of Atom See p 75 of text: electron energy levels are like rungs of a ladder. Ladder To climb to a higher level, you can’t put your foot at any level, you must place it on a rung Electron energy levels e-s must move to higher or lower e.l.’s in specific intervals

32 Electron Cloud Model of Atom
Electrons aren’t in perfect orbits. Energy levels are regions of space in which an e- is likely to be found most of the time. The area in which they move is like a cloud, an area of space surrounding the nucleus.

33 Electrons in Energy Levels
Atoms are arranged in energy levels (e.l.’s), at different distances from nucleus Close to nucleus = low energy Far from nucleus = high energy e-s in highest occupied level are “valence e-s” Only so many e-’s can fit in energy levels e-s fill lower e.l.’s before being located in higher e.l.’s* (* There are exceptions we will learn later!)

34 Electrons in Energy Levels
Only so many e-’s can fit in energy levels Energy Level # of electrons 1st 2 2nd 8 3rd * 4th *


36 Drawing Bohr Models Let’s practice drawing some atoms/ions In your teams, pick up enough of the following for your team: 1 white board per person 1 marker per person 1 paper towel per team (Please save a tree & share!)

37 Drawing Bohr Models Show # of protons and neutrons in the nucleus
Draw e.l.’s and show each electron in the proper e.l. Ex: Bohr Model of BORON-11

38 Practice Hydrogen-2 (Practice together) Helium-4 Lithium-6 Beryllium-8
Carbon-12 Magnesium-24

39 Lewis-Dot Diagrams Have 2 parts Chemical symbol of element
Valence e-s, represented by dots Are placed in one of four locations Above Below Right left Are not paired unless there is 1 e- in each location. Ex: Oxygen

40 Practice Lewis Dot Diagrams
TEACHER DEMONSTRATION Hydrogen Helium Lithium STUDENT PRACTICE Beryllium Boron Carbon

41 PRACTICE WORKSHEET Bohr Models Lewis dot diagrams

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