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Atoms, Molecules, and Ions Chapter 2 The Discovery of Atomic Structure The ancient Greeks were the first to postulate that matter consists of indivisible constituents. Later scientists realized that the atom consisted of charged entities. The Atomic Theory of Matter John Dalton: Each element is composed of atoms All atoms of an element are identical. In chemical reactions, the atoms are not changed. Compounds are formed when atoms of more than one element combine. Dalton’s law of multiple proportions: When two elements form different compounds, the mass ratio of the elements in one compound is related to the mass ratio in the other by a small whole number. The Discovery of Atomic Structure Cathode Rays and Electrons A cathode ray tube (CRT) is a hollow vessel with an electrode at either end. A high voltage is applied across the electrodes. Cathode Rays and Electrons The voltage causes negative particles to move from the negative electrode to the positive electrode. The path of the electrons can be altered by the presence of a magnetic field. The Discovery of Atomic Structure Cathode Rays and Electrons The Discovery of Atomic Structure Cathode Rays and Electrons Consider cathode rays leaving the positive electrode through a small hole. •If they interact with a magnetic field perpendicular to an applied electric field, the cathode rays can be deflected by different amounts. •The amount of deflection of the cathode rays depends on the applied magnetic and electric fields. •In turn, the amount of deflection also depends on the charge to mass ratio of the electron. The Discovery of Atomic Structure Cathode Rays and Electrons In 1897, Thomson determined the charge to mass ratio of an electron to be 1.76 108 C/g. Goal: find the charge on the electron to determine its mass. The Discovery of Atomic Structure Millikan Oil Drop Experiment The Discovery of Atomic Structure Cathode Rays and Electrons Consider the following experiment: •Oil drops are sprayed above a positively charged plate containing a small hole. •As the oil drops fall through the hole, they are given a negative charge. •Gravity forces the drops downward. The applied electric field forces the drops upward. •When a drop is perfectly balanced, the weight of the drop is equal to the electrostatic force of attraction between the drop and the positive plate. The Discovery of Atomic Structure Cathode Rays and Electrons Using this experiment, Millikan determined the charge on the electron to be 1.60 10-19 C. Knowing the charge to mass ratio, 1.76 108 C/g, Millikan calculated the mass of the electron: 9.10 10-28 g. With more accurate numbers, we get the mass of the electron to be 9.10939 10-28 g. The Discovery of Atomic Structure Radioactivity Consider the following experiment: •A radioactive substance is placed in a shield containing a small hole so that a beam of radiation is emitted from the hole. •The radiation is passed between two electrically charged plates and detected. •Three spots are noted on the detector: •a spot in the direction of the positive plate, •a spot which is not affected by the electric field, •a spot in the direction of the negative plate. The Discovery of Atomic Structure Radioactivity The Discovery of Atomic Structure Radioactivity A high deflection towards the positive plate corresponds to radiation which is negatively charged and of low mass. This is called b-radiation (consists of electrons). No deflection corresponds to neutral radiation. This is called g-radiation. Small deflection towards the negatively charged plate corresponds to high mass, positively charged radiation. This is called a-radiation (He atom). The Discovery of Atomic Structure The Nuclear Atom From the separation of radiation we conclude that the atom consists of neutral, positively, and negatively charged entities. Thomson assumed all these charged species were found in a sphere. The Discovery of Atomic Structure The Nuclear Atom Rutherford’s a-particle experiment: The Discovery of Atomic Structure The Nuclear Atom Rutherford carried out the following experiment: A source of a-particles was placed at the mouth of a circular detector. The a -particles were shot through a piece of gold foil. Most of the a-particles went straight through the foil without deflection. Some a-particles were deflected at high angles. If the Thomson model of the atom was correct, then Rutherford’s result was impossible. The Discovery of Atomic Structure The Nuclear Atom In order to get the majority of a-particles through a piece of foil to be undeflected, the majority of the atom must consist of a low mass, diffuse negative charge - the electron. To account for the small number of high deflections of the a-particles, the center or nucleus of the atom must consist of a dense positive charge. Atoms are mostly empty space!! The Discovery of Atomic Structure The Nuclear Atom Rutherford modified Thomson’s model as follows: assume the atom is spherical but the positive charge must be located at the center, with a diffuse negative charge surrounding it. The ModernView of Atomic Structure The atom consists of positive, negative, and neutral entities (protons, electrons, and neutrons). Protons and neutrons are located in the nucleus of the atom, which is small. Most of the mass of the atom is due to the nucleus. There can be a variable number of neutrons for the same number of protons. Isotopes have the same number of protons but different numbers of neutrons. Electrons are located outside of the nucleus. Most of the volume of the atom is due to electrons. The ModernView of Atomic Structure The ModernView of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers Atomic number (Z) = number of protons in the nucleus. Mass number (A) = total number of nucleons in the nucleus (i.e., protons and neutrons). A By convention, for element X, we write Z X Isotopes have the same Z but different A. The ModernView of Atomic Structure Isotopes, Atomic Numbers, and Mass Numbers Periodic Table The Periodic Table Columns in the periodic table are called groups (numbered from 1A to 8A or 1 to 18). Rows in the periodic table are called periods. Metals are located on the left hand side of the periodic table (most of the elements are metals). Non-metals are located in the top right hand side of the periodic table. Elements with properties similar to both metals and nonmetals are called metalloids and are located at the interface between the metals and non-metals. The Periodic Table The Periodic Table Some of the groups in the periodic table are given special names. These names indicate the similarities between group members: Group 1A: Alkali metals. Group 2A: Alkaline earth metals. Group 6A: Chalcogens. Group 7A: Halogens. Group 8A: Noble gases. Molecules and Molecular Compounds Molecules and Chemical Formulas •Molecules are assemblies of two or more atoms bonded together. •The chemical formula indicates •which atoms are found in the molecule, and in what proportion they are found. •Compounds formed from molecules are called molecular compounds. Molecules and Molecular Compounds Molecular and Empirical Formulas Molecular formulas give the actual numbers and types of atoms in a molecule. Examples: H2O, CO2, CO, CH4, H2O2, O2, O3, and C2H4. Empirical formulas give the relative numbers and types of atoms in a molecule. That is, they give the lowest whole number ratio of atoms in a molecule. Examples: H2O, CO2, CO, CH4, HO, CH2. Molecules and Molecular Compounds Picturing Molecules If the structural formula does show the shape of the molecule, then either a perspective drawing, ball-and-stick model, or space-filling model is used. Molecular Naming Two non-metals Same side of Table Naming 2 nonmetals Binary Molecular Compounds The most metallic element is usually written first (i.e., the one to the farthest left on the periodic table). Exception: NH3. If both elements are in the same group, the lower one is written first. Greek prefixes are used to indicate the number of atoms. Naming Inorganic Compounds Names and Formulas of Binary Molecular Compounds Ionic Compounds Opposite sides of Periodic Table MetalsNonmetals Ions and Ionic Compounds Ionic Compounds The majority of chemistry involves the transfer of electrons between species. Example: To form NaCl, the neutral sodium atom, Na, must lose an electron to become a cation: Na+. The electron cannot be lost entirely, so it is transferred to a chlorine atom, Cl, which then becomes an anion: Cl-. The Na+ and Cl- ions are attracted to form an ionic NaCl lattice which crystallizes. Ions and Ionic Compounds Important: note that there are no easily identified NaCl molecules in the ionic lattice. Therefore, we cannot use molecular formulas to describe ionic substances. Ions and Ionic Compounds When an atom or molecule loses electrons, it becomes positively charged. For example, when Na loses an electron it becomes Na+. Positively charged ions are called cations. Ions and Ionic Compounds When an atom or molecule gains electrons, it becomes negatively charged. For example when Cl gains an electron it becomes Cl-. Negatively charged ions are called anions. An atom / molecule can lose more than 1 electron. Naming Inorganic Compounds Names and Formulas of Ionic Compounds Name the cation then anion for the ionic compound. Example: BaBr2 = barium bromide. Ions and Ionic Compounds Predicting Ionic Charge The number of electrons an atom loses is related to its position on the periodic table. Metals tend to form cations whereas non-metals tend to form anions. Naming Inorganic Compounds Names and Formulas of Ionic Compounds Cations formed from a metal have the same name as the metal. Example: Na+ = sodium ion. If the metal can form more than one cation, then the charge is indicated in parentheses in the name (usually Transition Metals). Examples: Cu+ = copper(I); Cu2+ = copper(II). Naming Inorganic Compounds Names and Formulas of Ionic Compounds Monatomic anions (with only one atom) are called -ide. Example: Cl- is chloride ion. Exceptions: hydroxide (OH-), cyanide (CN-), peroxide (O22-). Polyatomic anions (with many atoms) containing oxygen end in -ate or -ite. (The one with more oxygen is called -ate.) Examples: NO3- is nitrate, NO2- is nitrite. Naming Inorganic Compounds Names and Formulas of Ionic Compounds Polyatomic anions containing oxygen with more than two members in the series are named as follows (in order of decreasing oxygen): per-…..-ate -ate -ite hypo-….-ite ClO4ClO3ClO2ClO - perchlorate chlorate chlorite hypochlorite Naming Inorganic Compounds Names and Formulas of Ionic Compounds Polyatomic anions containing oxygen with additional hydrogens are named by adding hydrogen or bi- (one H), dihydrogen (two H), etc., to the name as follows: CO32- is the carbonate anion HCO3- is the hydrogen carbonate (or bicarbonate) anion. H2PO4- is the dihydrogen phosphate anion. Naming Inorganic Compounds Names and Formulas of Acids The names of acids are related to the names of anions: -ide becomes hydro-….-ic acid; -ate becomes -ic acid; -ite becomes -ous acid. PERIODIC TRENDS General Periodic Trends Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Effective Nuclear Charge, Z* (kernel Charge) Z* is the nuclear charge experienced by the outermost electrons. Explains why Energy(2s) < Energy(2p) Z* increases across periodic table Orbitals in Many Electron Atoms Effective Nuclear Charge - Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge. - The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons. - As the average number of screening electrons (S) increases, the effective nuclear charge (zeff) decreases. zeff = z - S 48 Effective Nuclear Charge, Z* Atom Li B C N O F Z* Experienced by 2s Electrons in Valence Orbitals +1.28 +2.58 Increase in Z* +3.22 across a +3.85 period +4.49 +5.13 [Values calculated using Slater’s Rules] General Periodic Trends Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Atomic Size Size Increases going down a group. Because electrons are added further from the nucleus, there is less attraction. Size Decreases going across a period. Electrons are held closer. Atomic Size Size decreases across a period due to an increase in Z*. Each added electron feels a greater and greater (+) charge. Large Small Increase in Z* Atomic Radii Questions Of the elements magnesium, Mg, Chlorine, Cl, sodium Na, and phosphorus P which has the largest atomic radius? Explain. Of the elements Ca, Be, Ba, and Sr, which has the largest atomic radius? Explain. Of the elements Br, At, F, I and Cl which one has the smallest atomic radius and which has the largest atomic radius? General Periodic Trends Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Trends in Ionization Energy 1st Ionization energy (kJ/mol) 2500 He Ne 2000 Ar 1500 Kr 1000 500 0 1 H 3 Li 5 7 9 11 Na 13 15 17 19 K 21 23 25 27 29 31 Atomic Number 33 35 Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + e- Ionization Energy IE = energy required to remove an electron from an atom in the gas phase. Mg (g) + 738 kJ ---> Mg+ (g) + eCalled 1st IE Mg+ (g) + 1451 kJ ---> Mg2+ (g) + eCalled 2nd IE Ionization Energy See Screen 8.12 Mg (g) + 735 kJ ---> Mg+ (g) + eMg+ (g) + 1451 kJ ---> Mg2+ (g) + e- Mg2+ (g) + 7733 kJ ---> Mg3+ (g) + eEnergy cost is very high to dip into a shell of lower n. Does not happen. Trends in Ionization Energy Trends in Ionization Energy IE increases across a period because Z* increases. Metals lose electrons more easily than nonmetals. Metals are good reducing agents. Nonmetals lose electrons with difficulty. Trends in Ionization Energy IE decreases down a group Because size increases. Electrons are held less closely Questions Consider the four hypothetical main group elements Q, R ,T, X with the outer electron configurations indicated below. Then answer the questions that follow. Q = 3s23p5 R=3s1 T=4d105s25p5 X=4d105s25p1 a. b. c. d. Identify the block location of each hypothetical element. Which of these elements are in the same period? Which are in the same group? Which element would you expect to have the highest first IE? Which would have the lowest first IE? Which element is most likely to form a 1+ ion? General Periodic Trends Atomic and ionic size Ionization energy Electron affinity Higher effective nuclear charge Electrons held more tightly Larger orbitals. Electrons held less tightly. Electron Affinity New terms required Anion= negative charged atom Cation= positive charged atoms Easy to remember!!! Electron Affinity A few elements GAIN electrons to form anions. Halogens are the easiest. Electron affinity is the energy involved when an atom gains an electron to form an anion. A + e- ---> AE.A. = ∆E Electron Affinity Trends - numbers indicate an ease in adding electrons. Indicates how much energy is given off in the process. It is like a bank money (energy in, uses) = positive, spending money (energy out, releases) = negative numbers. Ex. EA of F =-339.9 kJ/mol (easy to do) EA of Mg = 0 kJ/mol (more difficult to do) Electron affinity Trends Cont In general electrons are more difficult to add going down a group. Because Nuclear charge increases Atomic radius increases down a group decreasing electron affinity Adding more electrons Is it possible to add more electrons to something that is already negatively charged. Yes. But it is more difficult. Only certain cases when it creates a noble gas configuration do electrons continue to add like in O-2 or N-3 Ion Radii Li,152 pm 3e and 3p Does+the size of the atoms and ions go + Li , 60 pm up or down when 2e and 3p losing an electron to form a cation? Ion Sizes + Li,152 pm 3e and 3p Li + , 78 pm 2e and 3 p Forming a cation. CATIONS are SMALLER than the atoms from which they come. The electron/proton attraction has gone UP and so size DECREASES. Atoms also decrease in size. Ion Sizes Does the size of the atoms and ions go up or down when gaining an electron to form an anion? Ion Sizes F, 71 pm 9e and 9p F- , 133 pm 10 e and 9 p Forming an anion. ANIONS are LARGER than the atoms from which they come. The electron/proton attraction has gone DOWN and so size INCREASES. Atom size also increases. Trends in ion sizes are the same as atom sizes. Trends in Ion Sizes Active Figure 8.15 Atomic Radii Size Trends Metals at the left end tend to form cations Nonmetals at the upper right tend to form anions. Electronegativity A measure of the ability of an atom in a chemical compound to attract electrons. Most electronegative is Fluorine. Least electronegative are alkali and alkaline metals Most electronegative are the halogens nitrogen and oxygen. Trends Electronegativty is strongest at the upper right of periodic table Weakest in the lower right. Decrease or stay the same down a group. Electronic Structure of Atoms Chapter 7 78 The Wave Nature of Light - Visible light is a small portion of the electromagnetic spectrum 79 The Wave Nature of Light Frequency (v, nu) – The number of times per second that one complete wavelength passes a given point. Wavelength (l, lambda) – The distance between identical points on successive waves. 80 lv=c c = speed of light, 2.997 x 108 m/s V=frequency (1/s = hertz= Hz) l =wavelenth (m) The Wave Nature of Light - We can also say that light energy is quantized - This is used to explain the light given-off by hot objects. - Max Plank theorized that energy released or absorbed by an atom is in the form of “chunks” of light (quanta). E=hv h = plank’s constant, 6.63 x 10-34J/s E= Energy in system (Joules= kg . m2/s2) V= frequency (Hz=1/s) - Energy must be in packets of (hv), 2(hv), 3(hv), etc. 81 Quantized Energy and Photons The Photoelectric Effect 82 Quantized Energy and Photons Einstein’s The Photoelectric Effect - The photoelectric effect provides evidence for the particle nature of light. - It also provides evidence for quantization. - If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal. - Below the threshold frequency, no electrons are ejected. - Above the threshold frequency, the number of electrons ejected depend on the intensity of the light. 83 Quantized Energy and Photons The Photoelectric Effect - Einstein assumed that light traveled in energy packets called photons. - The energy of one photon, E = hn.= hc/ l - This equation means that the energy of the photon is proportional to its frequency. 84 Chapter 6 De Broglie’s Contribution Thought: if waves show particle properties, can particles show wave properties? Answer is yes Mass=h / l x v (Einstein) so, Rearrrange l = h / mass x velocity Example Find the wavelength for a ball with a mass of .10 kg thrown at 35 m/s. l = h / mass x velocity Plug in numbers, 6.626 x .m/s 10-34 kg.m (.10 kg) ( 35 m/s) = 1.9 x 10-34 m Conclusions All matter exhibits both particulate and wave properties. Large mater exhibit mostly particulate properties ex. Baseballs, footballs, cannon balls, Rocks Small matter exhibit mostly wave properties ex. Photons, Intermediate matter exhibit like electrons demonstrate both. Quantum Model of the Atom Bohr’s Model of the Hydrogen Atom Bohr’s Model - Assumed that a single electron moves around the nucleus in a circular orbit. - The energy of a given electron is assumed to be restricted to a certain value which corresponds to a given orbit. 89 Bohr’s Model of the Hydrogen Atom Line Spectra for light is just like hydrogen 90 Chapter 6 Bohr’s Model of the Hydrogen Atom Bohr’s Model - Since the energy states are quantized, the light emitted from excited atoms must be quantized and appear as line spectra. 91 Chapter 6 Bohr’s Model of the Hydrogen Atom Line Spectra Line spectra can be “explained” by the following equation: -18 Z E - 2.179 x10 J 2 n n=energy level Z= number of protons 92 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Important Features - The first orbit in the Bohr model has n = 1 and is closest to the nucleus. - The furthest orbit in the Bohr model has n close to infinity and corresponds to zero energy. - Electrons in the Bohr model can only move between orbits by absorbing and emitting energy in quanta (hn). 93 Bohr’s Model of the Hydrogen Atom Bohr’s Model – Line Spectra Ground State – When an electron is in its lowest energy orbit. Excited State – When an electron gains energy from an outside source and moves to a higher energy orbit. 94 Chapter 6 More Equations ΔE= energy of final – energy of initial Will tell us if an atom has gained or lost energy. Negative sign means more stable and a loss of energy. Plus sign means less stable and a gain of energy. From that we can then determine the wavelength of that emitted photon using the equation: l hc/ ΔE Example Calculate the energy required to excite the hydrogen electron from level 1 to level 2. Use equation Z E - 2.179 x10 -18 J 2 n -2.179 x 10-18 (12/12) = -2.179 x 10-18 J For n=2 -2.179 x 10-18 (12/22) = -5.445 x 10-19 J For n=1 So, ΔE= energy of final – energy of initial ΔE= (-5.445 x 10-19 J) - (-2.179 x 10-18 J) 1.633 x 10-18 J (+ sign means energy is absorbed) 2 Things to Remember About the Bohr Model Only works with Hydrogen As you get closer to the ground state, energy is being released. This was a good start, and helps us To understand the atom, it was fundamentally Incorrect. The Wave Behavior of Matter - Remember DeBroglie proposed that there is a wave/particle duality. - Knowing that light has a particle nature, it seems reasonable to assume that matter has a wave nature. - DeBroglie proposed the following equation to describe the relationship: h l mv 98 The Wave Behavior of Matter The Uncertainty Principle Heisenberg’s Uncertainty Principle - on the mass scale of atomic particles, we cannot determine exactly the position, speed, and direction of motion simultaneously. - For electrons, we cannot determine their momentum and position simultaneously. 99 Quantum Mechanics - These theories (wave/particle duality and the uncertainty principle) mean that the Bohr model needs to be refined. Quantum Mechanics 100 Quantum Mechanics - The path of an electron can no longer be described exactly, now we use the wavefunction(y). Wavefunction (y) – A mathematical expression to describe the shape and energy of an electron in an orbit. Probability density = y2 101 Quantum Mechanics Quantum Numbers - The use of wavefunctions generates four quantum numbers. - Principal Quantum Number (n) - This is the same as Bohr’s n - Allowed values: 1, 2, 3, 4, … (integers) 102 Quantum Mechanics Quantum Numbers Secondary (Azimuthal) Quantum Number (l) - Allowed values: 0, 1, 2, 3, 4, . , (n – 1) (integers) - Each l represents an orbital type 103 l orbital 0 s 1 p 2 d 3 f Quantum Mechanics Quantum Numbers Magnetic Quantum Number (ml ). - This quantum number depends on l. - Allowed values: -l +l by integers. - Magnetic quantum number describes the orientation of the orbital in space. 104 l Orbital ml 0 s 0 1 p -1, 2 d -2, -1, 0, +1 0, +1, +2 Quantum Mechanics Quantum Numbers Spin Quantum Number (s) - Allowed values: -½ +½. - Electrons behave as if they are spinning about their own axis. - This spin can be either clockwise or counter clockwise. 105 Chapter 6 Quantum Mechanics Orbitals and Quantum Numbers 106 Representation of Orbitals The s Orbitals - All s-orbitals are spherical. - As n increases, the s-orbitals get larger. - As n increases, the number of nodes increase. - A node is a region in space where the probability of finding an electron is zero. 107 Representation of Orbitals The s Orbitals 108 Representation of Orbitals The p Orbitals - There are three p-orbitals px, py, and pz. (The letters correspond to allowed values of ml of -1, 0, and +1.) - The orbitals are dumbbell shaped. 109 Representation of Orbitals The p Orbitals 110 Representation of Orbitals The d and f Orbitals - There are 5 d- and 7 f-orbitals. - Four of the d-orbitals have four lobes each. - One d-orbital has two lobes and a collar. 111 Representation of Orbitals The d and f Orbitals 112 Electron configuration Electron configuration is the arrangement of electron in an atom. All atoms have different arrangements Like everything in nature, they want to have the lowest possible energy called ground states. 3 Rules Determining Ground State Electron Configurations Aufbau principle Pauli excusion principle Hund’s rule Aufbau Principle Electrons will occupy the lowest energy orbital that can receive it. Notice the order changes after 3p. It does not always go in order. Help is on the way!!! If you can count to seven then you can determine which energies are required. Or Use your periodic table Electron Configurations and the Periodic Table 117 Chapter 6 Pauli Exclusion Principle No two electrons can have the same 4 quantum numbers. This means that the numbers may look the same, but the spin will be different for electrons that occupy different orbitals. Ex. 1s2 2s2 2p3 1s2 2s2 2p6 3s2 3p6 Hund’s Rule Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must all have the same spin. Ex. OR 3 Methods to representing electron configurations Orbital notation Electron configuration notation Noble Gas notation Orbital Notation An unoccupied orbital is represented by a line with the orbital name under it. Electron spins are then entered. Ex. H or He 1s1 1s2 Electron configuration notation Eliminates the lines and arrows of orbital notation. The number of electrons in a sublevel is shown by adding a superscirpt to the sublevel designation. Ex. H 1s1 He 1s2 Be 1s2 2s2 Noble gas notation Same as electron configurations, but it uses noble gases to shorten the work. Noble gases are placed in brackets to indicate full electron shells and then the proper amount of energy levels is added after it. Ex. K [Ar] 4s1 Or Zn [Ar]3d10 4s2