Chapter 2 Atoms, Molecules, and Ions · 2017-05-16 · destroyed in chemical reactions.) Note: This is always true for chemical changes. However, we now know that atoms of one element

Chapter 2 Atoms, Molecules,

and Ions

Prepared by John N. Beauregard

Starting from a presentation by James F. Kirby

Quinnipiac University Hamden, CT

Lecture Presentation

© 2015 Pearson Education, Inc.

Atoms, Molecules, and Ions


Atomic Theory of Matter •  The idea that atoms are the

fundamental building blocks of matter was originated ~2400 years ago by the Greek philosopher Democritus.

•  It reemerged in the early 19th century

in a scientific theory championed by the British chemist John Dalton.

•  Dalton’s atomic theory explains

several scientific laws: the conservation of mass in chemical reactions, the constant composition of chemical compounds, and the law of multiple proportions.



Postulate #1 of Dalton’s Atomic Theory

Each element is composed of extremely small particles called atoms. (Or the atom is the fundamental building block of chemical matter.)

Note: As far as we know, this is absolutely true.




All atoms of a given element are identical to one another in mass and other properties, but the atoms of one element are different from the atoms of all other elements. Note: We now know this is not exactly true, as most elements have more than 1 isotope. However, all isotopes of a given element have identical reactivity, and the average atomic mass is the same for any sample of a given element. So, at the macroscopic level, we can treat the atoms of an element as though they have all the same mass.




Atoms of an element are not changed into atoms of a different element by chemical reactions (or atoms are neither created nor destroyed in chemical reactions.) Note: This is always true for chemical changes. However, we now know that atoms of one element do sometimes changes into atoms of other elements. However, this occurs in processes that fall within the realm of nuclear physics, not chemistry.



Postulate #4 of Dalton’s Atomic Theory Atoms of more than one element combine to form compounds; a given compound always has the same relative numbers and kind of atoms.

Click here to view a YouTube video tutorial on the atomic theory of matter. The video also shows how atomic theory explains the law of constant composition of chemical compounds. This law was introduced in Chapter 1.

https://www.youtube.com/watch?v=ml3QAV4V_a0



Law of Conservation of Mass

The total mass of substances present at the end of a chemical process is the same as the total mass of substances present before the process took place. Ø Dalton’s atomic theory explains this law by saying

that, during a chemical reaction, the atoms composing the starting chemicals (reactants) are simply re-arranged to form new substances (products). Thus, conservation of mass equates to conservation of atoms.



Law of Multiple Proportions If two elements, A and B, form more than one compound, the masses of B that combine with a given mass of A are in the ratio of small whole numbers.

Click here to view a YouTube video tutorial on the Law of Multiple Proportions. The law is explained in the context of the atomic theory of matter.

https://www.youtube.com/watch?v=l-IPd_r7ytw



Discovery of Subatomic Particles In Dalton’s original model, the atom was the smallest particle possible. However, subsequent experimental discoveries showed that the atom itself was made up of smaller particles. Ø Electrons (cathode ray and oil drop experiments) Ø Radioactivity (Note: an important topic but not covered

on the AP Exam.) Ø Nucleus (gold foil experiment): composed of protons

and neutrons.



Discovery of The Electron: J.J. Thomson’s Cathode Ray Experiment (1897)

English physicist J.J. Thomson found that streams of negatively charged particles (electrons) emanate from cathode tubes, causing fluorescence. (Note: the negative charge of the particles is evident from the direction of their deflected path.)



Click here to view a YouTube video of a cathode ray demonstration. The significance of Thomson’s observations is also discussed.

•  Thomson determined the electric field strength needed to straighten the path of a cathode ray deflected by a magnetic field of known strength. The allowed him to determine the charge/mass ratio of the electron: –1.76 x 108 coulombs/gram (C/g).

•  He obtained the same result no matter which metal was used for the

cathode. This indicated that the electron was common to all types of atoms.

https://www.youtube.com/watch?v=O9Goyscbazk



Millikan Oil-Drop Experiment: Determination of the Charge of an Electron

§  In 1909, Robert Millikan determined the charge on the electron to be 1.602 x 10–19 Coulombs (C).

§  He combined this with Thomson’s result for the charge/mass ratio of the electron and was able to calculate the mass of an electron.

Click here to view a YouTube video tutorial on Millikan’s experiment. The video also reviews Thomson’s CRT experiment.

https://www.youtube.com/watch?v=DnO3mkCLYsI



Sample Exercise: Given both Thomson’s experimental result for the charge to mass ratio for the electron (–1.76 x 108 C/g) and Millikan’s experimental result for the charge of the electron (–1.602 x 10–19 C), calculate the mass of a single electron. How does your result compare to the mass of a hydrogen atom (1.68 x 10–24 g)?



Radioactivity (Note: This material is not covered on the AP Exam.)

•  Radioactivity is the spontaneous emission of high-energy radiation by an atom.

•  It was first observed by Henri Becquerel (1896). Marie and Pierre Curie also made important contributions to our early understanding of this phenomenon.

•  Rutherford (prior to his gold foil experiment) showed that radioactivity was associated with the transmutation of one element into another. He also identified three types of radiation: alpha, beta, and gamma rays.

•  We now understand that radiation corresponds to either particles (alfa or beta) or photons (gamma) that are ejected from the nucleus of an atom as it changes from one element to another (i.e. during a “nuclear reaction”).



Types of Radioactivity (Note: This material is not covered on the AP Exam.)

Three types of radiation discovered by Ernest Rutherford: –  Alpha (α) particle: positively charged. Eventually it was found that it is

a helium nucleus emitted from the nucleus of a heavy atom as it decays into lighter atom.

–  Beta (β) particle: negatively charged. Eventually it was determined that it is an electron, but one emitted from a decaying nucleus (NOT an electron cloud).

–  Gamma (γ) ray: Uncharged. Eventually it was found to be a photon of high energy light emitted from a decaying nucleus.



Radiation is the Result of Nuclear Decay (Note: Not covered on the AP Exam)

Example of an Alpha Decay Nuclear Reaction High energy helium nucleus (2+ charge not shown) ejected from parent nucleus

These nuclear reactions involve the conversion of mass into energy, which manifests itself as high K.E. of the alpha and beta particles and high energy electromagnetic radiation (gamma rays)

Example of an Alpha Decay Nuclear Reaction High energy electron (–1 charge depicted as a superscript) ejected from parent nucleus, came from a proton that changed to a neutron



In Summary: •  Nuclear Chemistry (Chapter 21 in Brown &

Lemay) is not covered on the AP Exam. So you won’t be tested on it in this class.

•  Neither will you have to write nuclear reactions. •  Note: It may be helpful to have some familiarity

with alpha radiation due to its importance in Rutherford’s gold foil experiment, which is covered later in this presentation.



The Plum Pudding Model: Early Attempt to Explain how the Atom was Constructed

•  Around 1900 J.J. Thomson proposed his “plum pudding” model the atom.

•  It pictured the net neutral

atom as a sphere of positive matter with negative electrons embedded in it.



Discovery of the Nucleus by Ernest Rutherford

Around 1910 Ernest Rutherford shot α particles at a thin sheet of gold foil and observed the pattern of scatter of the particles.

Click here to view a video tutorial on Rutherford’s experiment and how his results led Rutherford to propose the nuclear model of the atom to replace of the “plum pudding” model.

https://www.youtube.com/watch?v=dNp-vP17asI



Description of Rutherford’s Results Ø  Most alpha particles either passed

straight through the gold foil or were only slightly deflected

Ø  However, around 1 in 20,000 alpha particles were scattered at large angles form the foil.

Ø  Since even a few particles were deflected at large angles, Thomson’s “plum pudding” model could not be correct.

(What had Rutherford expected based on the “plum pudding” model and why?)



The Nuclear Model of the Atom •  Rutherford postulated a very small,

dense nucleus with the electrons surrounding it.

•  Most of the volume is empty space. •  Atoms are very small: 1 – 5 Å (or

100 – 500 pm). •  Later other subatomic particles

(protons and neutrons) were discovered as parts of the nucleus.

Click here to view a demonstration of Rutherford’s experiment using detectors based on modern technology. The video also covers Rutherford’s original experiment and how his results led to the nuclear model of the atom.

https://www.youtube.com/watch?v=XBqHkraf8iE



Summary of Subatomic Particles •  Protons (+1) and electrons (–1) have equal but opposite charges.

Neutrons are have no charge. •  Protons and neutrons have essentially the same mass (about 1 amu

each). The mass of an electron (about 1/1800 amu) is negligible in comparison.

•  Protons and neutrons, which carry most of the mass, are found in the nucleus. Protons give the nucleus a positive charge.

•  Electrons travel around the nucleus and account for all of the negative charge and almost al of the volume of the atom.



The Atomic Mass Unit (amu) •  Atoms have extremely small masses; the heaviest

known atoms have a mass of only 4 × 10–22 g. •  Atomic Mass Unit (amu):

–  a convenient unit for expressing the masses of atoms and molecules.

–  By definition, 1 amu = 1/12 mass of a C-12 atom. (Note: A C-12 atom contains 12 protons, 12 neutrons, and 12 electrons.)



Isotopes •  atomic number (Z): the number of protons in the nucleus of an atom. Each

atom of a given element contains the same number of protons (or same atomic number). For instance, all carbon atoms have Z = 6.

•  Isotopes: atoms of the same element with different masses but identical chemical reactivity. Each isotope of a given element has the same number of protons (and electrons) but a different number of neutrons (N).

•  mass number (A): defined A = Z + N. Both protons and neutrons have masses of about 1 amu; the electron mass is negligible in comparison. Thus, the mass number gives the approximate the mass of an atom in amu. For example, the mass of a carbon-14 atom is approximately 14 amu.



Isotopic Symbols of Elements Shown below is the isotopic symbol for the isotope carbon–12:

Ø  The atomic number (Z) is written as a subscript BEFORE the element symbol.

Ø  The mass number (A) is is written as a superscript BEFORE the symbol. Ø  For charged atoms (ions), the charge is written as a superscript AFTER

the symbol. (See the following video for examples.)

Click here to view a tutorial video on writing isotopic symbols. The video covers examples for both neutral atoms and ions.

https://www.youtube.com/watch?v=AYDLg5B18dY&feature=relmfu



Average Atomic Mass (or Atomic Weight) •  Each sample of a given element contains exactly the same isotopes in the

same relative abundances. •  In the real world we use large amounts of atoms and molecules. So we

use average masses in calculations, since the average atomic mass is the same for any sample of a given element.

•  An average mass (or atomic weight) is found using all isotopes of an element weighted by their relative abundances:

Atomic Weight = Ʃ [(isotope mass) × (fractional natural abundance)]

Click here to view a tutorial on calculating average atomic masses.

Click here to view a second tutorial with additional examples of average atomic mass calculations.

https://www.youtube.com/watch?v=G6A4UQX7QRw

https://www.youtube.com/watch?v=b2W9uZrtavE



Mass Spectrometry: Experimental Technique Used to Obtain Atomic Mass Data

•  Common experimental technique for determining the masses of atoms and molecules.

•  We will focus on how it is used to determine both the number of isotopes for a given element and the precise mass of the atoms of each isotope.

•  The schematic below depicts a mass spectroscopy experiment on the element chlorine, which has two isotopes: 35Cl and 37Cl.

Schematic Diagram of a Mass Spectrometer



The Mass Spectrum of Atomic Chlorine Obtained from the Experiment Depicted on the Previous Slide

Click here to view a tutorial on mass spectrometry. The video shows how to analyze the mass spectrum of an element and use the data to determine the average atomic mass of that element.

Provides the following information: 1.  Precise mass of each isotope

of the element (relative to that of carbon 12).

2.  Abundance of each isotope of relative to the other isotopes of the same element.

https://www.youtube.com/watch?v=zYgMxINkgLU



Sample Problem: Do the following calculations based on the mass spectrum for chlorine shown on the previous slide.

1) Determine the %-abundance of each isotope of chlorine 2)  Find the (average) atomic mass of chlorine



The Periodic Table •  a systematic organization of the elements: in order of atomic number. •  the atomic number of an element is at the TOP of its corresponding box. •  the atomic weight of an element usually appears at the BOTTOM of the box.

(although not shown on this version of the Periodic Table.)



The Periodic Law When the elements are arranged in order of

increasing atomic number, a repeating pattern is observed in their properties.



Periodic Table: Periods and Groups

Periods §  Horizontal rows on the

periodic table §  Each corresponds to

one full cycle of the repeating properties of the elements

Groups (or Families) o  Vertical columns in the

periodic table o  Elements in the same

group have similar chemical properties.



These Five Groups Are Known by Special Names:



Periodic Table: Metals, Nonmetals, and Metalloids

Metals •  Located on the left side

of the periodic table. •  Metallic properties

include Ø  shiny luster. Ø  conductive of heat and

electricity. Ø Malleable and ductile Ø Solid state at room

temperature (except mercury).




Nonmetals •  located on the right side

of the periodic table (except for H).

•  Nonmetallic properties include: Ø  Poor conductors of heat and

electricity Ø  can be solid (e.g. C), liquid

(e.g. Br), or gas (e.g. Ne) at room temperature.

Ø  Brittle when solid




Metalloids •  Located on the step-

like strip dividing the metals and nonmetals (except Al, Po, and At).

•  Their properties are sometimes like metals and sometimes like nonmetals.



Chemical Formulas •  The subscript to the right of an

element symbol tells the number of atoms of that element in one molecule (or formula unit) of the substance.

•  Molecular substances: Ø  composed of neutral molecules

and typically contain only nonmetals (or sometimes a nonmetal and a metalloid).

Ø chemical formula gives the number of each type of atom contained in one molecule of the substance.



Molecular Elements Some elements can exist in molecular form (i.e. molecules with just one type of atom)

Ø  P4 is a molecular form of phosphorus Ø  S8 is a molecular form of sulfur Ø  C60 is a molecular form of carbon Ø  The most stable form of each of the following

seven elements is a diatomic molecule o  Hydrogen (H2) o  Nitrogen (N2) o  Oxygen (O2) o  Fluorine (F2) o  Chlorine(Cl2) o  Bromine(Br2) o  Iodine(I2)



Types of Formulas •  Empirical formula: gives the lowest whole-

number ratio of atoms of each element in a compound.

•  Molecular formula: gives the exact number of

atoms of each element in a compound. •  If we know the molecular formula of a compound,

we can determine its empirical formula. (This is covered in Chapter 3.) The converse is not true!



Types of Formulas •  Structural formula: shows the

order in which atoms in a molecule are attached but does NOT depict the three-dimensional shape of molecules.

•  Perspective (aka wedge and stick) drawings: show both the order of attachment of the atoms and their three-dimensional arrangement in space. This can also be demonstrated using molecular models.



Ions: Charged Particles •  Monatomic Ion: a single atom with a net charge. •  Cations: formed when one or more electrons is lost. Metals form

monatomic cations. •  Anions: formed when one or more electron is gained.

Nonmetals form monatomic anions. •  Note: The charges on the ions of main group elements (groups

1A–8A) can be predicted based on their group number.



Some Common Cations

Recall: Sometimes a Roman numeral is used to specify the charge of a metal cation. This is done whenever the metal for that cation can form cations with more than one possible charge.



Some Common Anions



Ionic Compounds v  Generally formed by reacting metals with nonmetals (e.g. NaCl) v  Electrons are transferred from the metal to the nonmetal. The

oppositely charged ions attract each other and form a crystal lattice. (Does not exist as individual molecules.)

v  Can also be formed with polyatomic ions (e.g. NH4Cl) v  Chemical formula gives the lowest whole-number combination of

cations an anions that is net neutral. (i.e. one formula unit of the compound).



Ionic Compounds: Simple Approach for Determining Formulas

Ø  First, determine the respective charges of the cation and anion. Ø  Then, since compounds are electrically neutral, one can

determine the formula of a compound this way: •  The charge on the cation becomes the subscript on the anion. •  The charge on the anion becomes the subscript on the cation. •  If these subscripts are not in the lowest whole-number ratio,

divide them by the greatest common factor.

Click here to view a tutorial on writing formulas for ionic compounds. Several examples are given.

https://www.youtube.com/watch?v=YZKNI907dmE



Oxyanion: polyatomic anion with one or more oxygen atoms around a central atom

•  Central atoms on the second row have a bond to, at most, three oxygens; those on the third row contain up to four.

•  Charges increase from right to left.



Patterns in Oxyanion Nomenclature •  When there are two oxyanions involving the same

element –  the one with fewer oxygens ends in -ite. –  the one with more oxygens ends in -ate.

•  Example 1: nitrite (NO2−) and nitrate (NO3

−) •  Example 2: sulfite (SO3

2−) and sulfate (SO42−)



Patterns in Oxyanion Nomenclature: What if the Central Atom Has More Than Two

Possible Oxyanions?

•  The one with the second fewest oxygens ends in -ite: ClO2− is

chlorite. •  The one with the second most oxygens ends in -ate: ClO3

− is chlorate.

•  The one with the fewest oxygens has the prefix hypo- and ends in -ite: ClO− is hypochlorite.

•  The one with the most oxygens has the prefix per- and ends in -ate: ClO4

− is perchlorate.



Nomenclature for Ionic Compounds The name for any ionic compound should be written as follows •  Give the cation name followed by the anion name (separated

by a space). •  If the cation can have more than one possible charge, write

the charge as a Roman numeral in parentheses. •  If the anion is an element, change its ending to –ide. If the

anion is a polyatomic ion, simply write the name of the polyatomic ion.

Click here to view a tutorial on naming ionic compounds. Lots of examples are given.

https://www.youtube.com/watch?v=CkCzceecCrc&amp;feature=relmfu



Acid Nomenclature 1.  If the anion in the acid ends in

-ide, change the ending to -ic acid and add the prefix hydro-.

Ø HCl: hydrochloric acid Ø HBr: hydrobromic acid Ø HI: hydroiodic acid

2.  If the anion ends in -ite, change the ending to -ous acid.

Ø HClO: hypochlorous acid Ø HClO2: chlorous acid

3.  If the anion ends in -ate, change the ending to -ic acid.

Ø HClO3: chloric acid Ø HClO4: perchloric acid

Click here to view a video tutorial on writing the names and formulas of acids. Several examples are given.

https://www.youtube.com/watch?v=0kQmZfKEWT8



Nomenclature of Binary Molecular Compounds

•  The name of the element farther to the left in the periodic table (closer to the metals) or lower in the same group is usually written first.

•  A prefix is used to denote the number of atoms of each element in the compound (mono- is not used on the first element listed, however).



Examples: Naming Binary Molecular Compounds •  The ending on the second element is changed to -ide.

Ø CO2: carbon dioxide Ø CCl4: carbon tetrachloride

•  If the prefix ends with a or o and the name of the element begins with a vowel, the two successive vowels are often elided into one.

Ø N2O5: dinitrogen pentoxide

Click here to view a tutorial video on writing the names of binary molecular compounds. Lots of examples are given.

Click here to view a second tutorial on writing the names and formulas of ionic compound. Many more examples are covered.

https://www.youtube.com/watch?v=NJn76CR70oU&feature=endscreen&NR=1

https://www.youtube.com/watch?v=H2qfn6AZLAc



Nomenclature of Organic Compounds (Note: Not Covered on the AP Exam)

•  Organic chemistry: the chemistry of carbon-based compounds.

•  Organic chemistry has its own system of nomenclature. •  Alkanes: the simplest hydrocarbons (compounds containing

only carbon and hydrogen and all single bonds. •  The first part of the names just listed correspond to the

number of carbons (meth- = 1, eth- = 2, prop- = 3, etc.).



Nomenclature of Organic Compounds (Note: Not Covered on the AP Exam)

•  When a hydrogen in an alkane is replaced with something else (a functional group, like -OH in the compounds above), the name is derived from the name of the alkane.

•  The ending denotes the type of compound. –  An alcohol ends in -ol.

Documents

Chapter 2 Atoms, Molecules, and Ions · 2017-05-16 · destroyed in chemical reactions.) Note: This is always true for chemical changes. However, we now know that atoms of one element