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Chapter 4 Chemical Foundations: Elements, Atoms, and Ions.

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Presentation on theme: "Chapter 4 Chemical Foundations: Elements, Atoms, and Ions."— Presentation transcript:

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2 Chapter 4 Chemical Foundations: Elements, Atoms, and Ions

3 Atomic Structure

4 Objectives Explain early models of the atom Differentiate between three types of subatomic particles Define isotope and calculate the atomic mass

5 Sections 4.1 Defining the Atom 4.2 Structure of the Nuclear Atom 4.3 Distinguishing Between Atoms

6 4.1 Defining the Atom Vocabulary - Atom - Dalton’s atomic theory Atom– is the smallest particle of an element that retains its identity in a chemical reaction. Dalton’s atomic theory- the first theory to relate chemical changes to events at the atomic level

7 Experiencing Atoms Atoms are incredibly small, yet they compose everything Atoms are the pieces of elements Properties of the atoms determine the properties of the elements If every atom within a pebble were the size of the pebble itself, then the pebble would be larger than Mt. Everest (~29,000 ft)

8 Early Models of the Atom Democritus  The Greek philosopher Democritus (460 BC –370 BC) was among the first to suggest the existence of atoms.  He believed that atoms were indivisible and indestructible

9 Early Models of the Atom John Dalton ( )  Dalton’s Atomic Theory - By using experimental methods, Dalton transformed Democritus’s ideas on atoms into a scientific theory.

10 Dalton’s Atomic Theory 1. All elements are composed of tiny indivisible particles called atoms. 2. Atoms of the same element are identical. The atoms of any one element are different from those of any other element.

11 3. Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds. 4. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element are never changed into atoms of another element in a chemical reaction.

12 What instruments are used to observe individual atoms?  Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes. Sizing up the Atom

13 4.2 Structure of the Nuclear Atom Electron– a negatively charged subatomic particle Cathode ray- a stream of electrons produced at the negative electrode (cathode) of a tube containing a gas at low pressure Proton– a positively charged subatomic particle found in the nucleus of an atom Neutron – a subatomic particle with no charge and a mass of 1 amu Nucleus- the tiny, dense central portion of an atom, composed of protons and neutrons Vocabulary - Electron - Cathode ray - Proton - Neutron - Nucleus

14 Subatomic Particles 4.2

15 Electrons English physicist J.J. Thomson discovered the electron (1897) Electrons are negatively charged subatomic particles. Cathode ray tube  Thomson performed experiments that involved passing electric current through gases at low pressure.

16 Electrons The U.S. physicist Robert A. Millikan (1868–1953) carried out experiments to find the quantity of an electron’s charge.  Using this charge and Thomson’s charge-to-mass ratio of an electron, Millikan calculated an electron’s mass.  An electron has one unit of negative charge, and its mass is 1/1840 the mass of a hydrogen atom.

17 Protons Eugen Goldstein (1850–1930)  Observed a cathode-ray tube and found rays traveling in the direction opposite to that of the cathode rays. - He concluded that they were composed of positive particles. - Such positively charged subatomic particles are called protons.

18 Neutrons Physicist James Chadwick (1891–1974)  Confirmed the existence of another subatomic particle: the neutron.  Neutrons are subatomic particles with no charge but with a mass nearly equal to that of a proton.

19 Subatomic Particles Table 4.1 summarizes the properties of electrons, protons, and neutrons. 4.2

20 Atomic Nucleus J.J. Thompson and others supposed the atom was filled with positively charged material and the electrons were evenly distributed throughout.  Plum pudding model Plum pudding model - This model of the atom turned out to be short-lived, however, due to the work of Ernest Rutherford (1871–1937).

21 The Atomic Nucleus Ernest Rutherford’s Portrait 4.2

22 Ernest Rutherford Discovered the nucleus (1911) Stated protons were inside of the nucleus Gold-Foil Experiment  Radioactive particles “shot” through gold foil - A majority of particles passed straight through foil - A small fraction of particles bounced off gold foil at large angles or bounced straight back

23 The Atomic Nucleus Rutherford’s Gold-Foil Experiment 4.2

24 The Atomic Nucleus 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. 4.2

25 4.3 Distinguishing Among Atoms Vocabulary - Atomic number - Mass number - Isotopes - Atomic Mass Unit - Atomic mass - Periodic table - Period - Group Atomic number– the number of protons in the nucleus of an atom of an element Mass number- the total number of protons and neutrons in the nucleus of an atom Isotopes- atoms of the same element that have the same atomic number but different atomic masses due to a different number of neutrons Atomic mass unit- a unit of mass equal to one-twelfth the mass of a carbon-12 atom Atomic mass- the weighted average of the masses of the isotopes of an element Periodic table- an arrangement of elements in which the elements are separated into groups based on a set of repeating properties Period- a horizontal row of elements in the periodic table Group- a vertical column of elements in the periodic table; the constituent elements of a group have similar chemical and physical properties

26 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Just as there are many types of dogs, atoms come in different varieties too. How can there be different varieties of atoms?

27 Atomic Number *Elements are different because they contain different numbers of protons. An element’s atomic number is the number of protons in the nucleus of an atom of that element. The atomic number identifies an element.

28 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. For each element listed in the table below, the number of protons equals the number of electrons. (Protons = Electrons) Interpret Data Atoms of the First Ten Elements NameSymbolAtomic number ProtonsNeutronsMass number Electrons HydrogenH11011 HeliumHe22242 LithiumLi33473 BerylliumBe44594 BoronB CarbonC NitrogenN OxygenO FluorineF NeonNe

29 Understanding Atomic Number The element nitrogen (N) has an atomic number of 7. How many protons and electrons are in a neutral nitrogen atom? How many protons and electrons are in a Krypton atom?

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

31 Mass Number Au is the chemical symbol for gold. 4.3 The atomic number is the subscript. The mass number is the superscript.

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33 Symbols Find the - number of protons - number of neutrons - number of electrons - Atomic number - Mass Number - Name Na 24 11

34 4.3 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 also have different mass numbers.

35 Isotopes 2 isotopes of Carbon: - C – 12 - C -14

36 Isotopes Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons. 4.3

37 Isotopes For example, hydrogen has 3 isotopes: Note that the correct way to write an isotopes is to write the name, followed by the mass number. IsotopeProtonNeutronElectron Hydrogen Hydrogen Hydrogen

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39 Isotopes & Their Uses The tritium content of ground water is used to discover the source of the water, for example, in municipal water or the source of the steam from a volcano.

40 Isotopes & Their Uses Bone scans with radioactive technetium-99.

41 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Why are atoms with different numbers of neutrons still considered to be the same element? Despite differences in the number of neutrons, isotopes are chemically alike. They have identical numbers of protons and electrons, which determine chemical behavior.

42 Learning Check An atom has 14 protons and 20 neutrons. A.Its atomic number is 1) 142) 163) 34 B. Its mass number is 1) 142) 163) 34 C. The element is 1) Si2) Ca3) Se D.Another isotope of this element is 1) 34 X 2) 34 X 3) 36 X

43 Atomic Mass How do you calculate the atomic mass of an element? 4.3

44 Atomic Mass An atomic mass unit (amu) is defined as one twelfth of the mass of a carbon-12 atom.  It is more useful to compare atoms using a reference isotope The mass listed in the periodic table is the average atomic mass  The atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample of the element. 4.3

45 Atomic Mass To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. 4.3

46 Atomic Mass Carbon has two stable isotopes: carbon-12, which has a natural abundance of percent, and carbon-13, which has a natural abundance of 1.11 percent. The mass of carbon-12 is amu; the mass of carbon-13 is amu. The atomic mass of carbon is calculated as follows: Atomic mass of carbon = ( amu x ) amu x ) = ( amu) + (0.144 amu) = amu

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48 Copyright © Pearson Education, Inc., or its affiliats. All Rights Reserved. The mass each isotope contributes to the element’s atomic mass can be calculated by multiplying the isotope’s mass by its relative abundance. The atomic mass of the element is the sum of these products. Analyze List the knowns and the unknown. 1 atomic mass of X = ? KNOWNS UNKNOWN Isotope 10 X: mass = amu relative abundance = 19.91% = Isotope 11 X: mass = amu relative abundance = 80.09% =

49 Use the atomic mass and the decimal form of the percent abundance to find the mass contributed by each isotope. for 10 X: amu x = amu for 11 X: amu x = amu Calculate Solve for the unknowns. 2

50 Add the atomic mass contributions for all the isotopes. Calculate Solve for the unknowns. 2 For element X, atomic mass = amu amu = amu

51 Evaluate Does the result make sense? 3 The calculated value is closer to the mass of the more abundant isotope, as would be expected.

52 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. Key Concepts Elements are different because they contain different numbers of protons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. To calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products.

53 Periodic Table-A Preview

54 Periodic Table A periodic table is an arrangement of elements in which the elements are separated into groups based on a set of repeating properties  It allows you to compare the properties of one element to another element. Period- a horizontal row of elements in the periodic table Group- a vertical column of elements in the periodic table; the constituent elements of a group have similar chemical and physical properties *Elements are listed in order of increasing atomic number, from left to right and top to bottom.

55 ATOMIC COMPOSITION Protons (p + ) Protons (p + ) + electrical charge + electrical charge mass = x g mass = x g relative mass = atomic mass units (amu) but we can round to 1 relative mass = atomic mass units (amu) but we can round to 1 Electrons (e - ) Electrons (e - ) negative electrical charge negative electrical charge relative mass = amu but we can round to 0 relative mass = amu but we can round to 0 Neutrons (n o ) Neutrons (n o ) no electrical charge no electrical charge mass = amu but we can round to 1 mass = amu but we can round to 1

56 IONS IONS are atoms or groups of atoms with a positive or negative charge. IONS are atoms or groups of atoms with a positive or negative charge. Taking away an electron from an atom gives a CATION with a positive charge Taking away an electron from an atom gives a CATION with a positive charge Adding an electron to an atom gives an ANION with a negative charge. Adding an electron to an atom gives an ANION with a negative charge. To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples: Na + Ca +2 I - O -2 To tell the difference between an atom and an ion, look to see if there is a charge in the superscript! Examples: Na + Ca +2 I - O -2 Na Ca I O Na Ca I O

57 PREDICTING ION CHARGES In general metals (Mg) lose electrons ---> cations metals (Mg) lose electrons ---> cations nonmetals (F) gain electrons ---> anions nonmetals (F) gain electrons ---> anions

58 Learning Check – Counting State the number of protons, neutrons, and electrons in each of these ions. 39 K + 16 O -241 Ca #p + ___________________ #n o ___________________ #e - ___________________

59 One Last Learning Check Write the nuclear symbol form for the following atoms or ions: A. 8 p +, 8 n, 8 e - ___________ B.17p +, 20n, 17e - ___________ C. 47p +, 60 n, 46 e - ___________

60 Charges on Common Ions By losing or gaining e-, atom has same number of e-’s as nearest Group 8A atom.

61 Regions of the Periodic Table

62 Group 1A: Alkali Metals Cutting sodium metal Reaction of potassium + H 2 O

63 Magnesium Magnesium oxide Group 2A: Alkaline Earth Metals

64 Group 7A: The Halogens (salt makers) F, Cl, Br, I, At


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