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UNIT 3 – ATOMIC THEORY CH1030 Mark Stacey ATOMIC THEORY A molecule is the smallest piece of a substance that retains the properties of that substance. For example, a water molecule (H2O) acts like water but if it is broken down further (into lone H or O) it no longer behaves as water. An atom is the smallest piece of a single element that still retains the properties of that element. A molecule can contain one or many atoms. Atoms are effectively the smallest whole units of matter. While atoms themselves are made of smaller parts, these parts do not exist independently in most everyday life situations. ATOMIC THEORY Atoms were first theorized by the ancient Greeks as a term to describe the smallest piece of a substance that still acted as that substance. However, this was long, long before an individual atom was very observed. This is because atoms are extremely tiny, in the range of 0.0000000001m in diameter. This means that even a few milligrams of a substance contains many, many atoms. The periodic table shows the ~120 types of atoms we have observed or predicted to exist. All matter is made up of one or a combination of these atoms. ATOMIC THEORY We describe different substances using chemical formulas, a shorthand method to describe the types and quantities of atoms in a single molecule. All types of atoms have one or two letter shorthand. The first letter is always capitalized and the second (if present) always is lowercase. Some elements are named based on older or non-English names and so their shorthand name may not align well with their modern English full name. Oxygen – O Carbon – C Sodium – Na (from Latin: natrium) Iron – Fe (from Latin: ferrum) ATOMIC THEORY An element is a chemical containing only one type of atom. Elements may exist as single atoms, or a grouped molecules – the important thing is that all the atoms are the same type. Compounds contain two or more types of atoms within their individual molecules. Simply having two types of atoms floating around separately is not a compound – that is a mixture. Mixtures contain two or more types of molecules. These molecules themselves can be elements or compounds. ATOMIC THEORY ATOMIC THEORY (Pure) Element (Pure) Compound Mixture Types of Atoms 1 2+ 2+ Types of Molecules 1 1 2+ Number of Atoms per Molecule 1+ 2+ n/a SUB-ATOMIC PARTICLES Atoms may be the smallest functional units of matter, but they are not the smallest things to exist. Atoms are made of three major subatomic particles: Neutrons (n) make up the core of an atom called a nucleus. These particles have no charge to them and are largely only contribute to the size and mass of an atom. Protons (p+) are also found within the nucleus. These have a similar size and mass to neutrons, but also have a positive charge. This makes the nucleus as a whole have a net positive charge. SUB-ATOMIC PARTICLES Electrons (e-) do not exist within the nucleus of the atom. Instead, they orbit the nucleus from a relatively far distance. Electrons are much, much smaller and lighter than protons or neutrons. Electrons have negative charge, equal and opposite to a proton’s positive charge. Electrons are kept in their orbits partially by this positive-negative attraction. A representation of an atom using the “planetary model”. Note that the sizes and distances shown here are not realistic. Electrons are significantly smaller than protons or neutrons, and the distance from electrons to the nucleus is significantly larger. SUB-ATOMIC PARTICLES As electrons are more “accessible” to the outside, it is electrons that are more readily lost, gained or shared to other atoms. Most chemical changes involve the redistribution of electrons between atoms. The nucleus is significantly more stable. However, nuclei can be joined together or split apart. These types of reactions take tremendously more effort to accomplish compared to regular chemical reactions involving electrons. As these reactions involve the nucleus, these are called nuclear reactions. SUB-ATOMIC PARTICLES Sub-Atomic Particle Neutron Location in Atom Nucleus Charge None Size 10-15 m Mass Reactivity 10-24 g Nuclear Reactions Only Proton Nucleus Positive 10-15 m 10-24 g Nuclear Reactions Only Electron Orbitals Around Nucleus Negative 10-16m 10-28 g Any Chemical Reaction MODELS OF THE ATOM While the concept of the atom has existed for two thousand years, the actual makeup of atoms themselves has only been explored in the past hundred years. The Greek idea of the atom as a single, whole, indivisible particle was codified into what is called the Dalton model of the atom. Under the Dalton model: The smallest pieces of an element possible are atoms Every element has a different atom unique to it Atoms cannot be further divided Atoms are the smallest particles to exist MODELS OF THE ATOM At the end of the 19th century, the electron was discovered to exist. As well, the electron was observed to be something that could be isolated from all elements, meaning that all atoms had to be partially made up of these electrons. Electrons were discovered to be negative in charge. As atoms are neutral in charge, scientists deduced that there must also be a positively-charged species in the atom to balance out the electron. The proton was deduced even before it was isolated. This created the Thompson model: Atoms are no longer singular pieces, but instead contain protons and electrons (and potentially other yet-to-be-discovered parts Electrons and protons are non-specifically arranged in the “mixture” that is an atom Thompson likened the atom to a plum pudding, were electron and protons are suspended in the atom, like the plums in the pudding. MODELS OF THE ATOM MODELS OF THE ATOM Approximately 20 years later, experiments were done to prove the existence of the nucleus. This was done by sending a beam of radiation through an atom that would only react to protons and observing the “shadow” cast from the protons. If Thompson’s model were true, the “shadow” would appear as a somewhat random distribution, as the protons should have no organization. Instead, it was observed that the majority of the atom cast no “shadow” at all, and only a very small area at the centre did. This lead to the Rutherford model: Protons are densely-packed into a central nucleus Protons only exist within this nucleus. MODELS OF THE ATOM MODELS OF THE ATOM Not long after, the neutron was discovered to exist. It’s lack of a charge made it harder to observe than the charged particles. The neutron was discovered to also exist with in the nucleus along with the proton. As time went on, the Bohr model would be developed, which explained that electrons not only orbited the nucleus randomly, but in a very specific series of orbits. This is sometimes known as the planetary model, as it resembles a star and its orbiting planets. Recently, the idea of electrons forming discrete circular orbits has been replaced with the idea that electrons exist within an orbital cloud. This is due to the fact that objects as small as an electron can behave differently than larger objects. This is the quantum model. MODELS OF THE ATOM ATOMS AND ELEMENTS All atoms are a combination of protons, neutrons and electrons. What makes one atom different from another (and therefore what makes one element different from another) is the numbers of protons, neutrons and electrons. As protons make up the nucleus and have a charge, we define each element by the number of protons. Neutrons also make up the nucleus, but as they are neutrally-charged, they do not play into the behaviors of the atom significantly. The number of protons in an atom is known as the atomic number (Z). The simplest element, hydrogen has one proton (Z=1), while the heaviest elements currently known have over 110 protons (Z=110+) ATOMS AND ELEMENTS The atomic number is the large number displayed with an element on the periodic table. Atomic numbers are always whole numbers (no decimals), as there cannot be a fraction of a proton. Atomic numbers cannot be less than 1. All atoms have at least one proton. ATOMS AND ELEMENTS The number of protons defines which element an atom is. For example, an atom with 8 protons is oxygen, regardless of the number of neutrons and electrons. If there are 8 protons, then it is some “variant” on oxygen. As neutrons have no charge, they effectively only contribute mass and size to the nucleus. Therefore if a oxygen atom with 10 neutrons is not terribly different from an oxygen with 8 neutrons, other than being heavier. ATOMS AND ELEMENTS The various versions of an atom containing differing numbers of neutrons are called isotopes. Isotopes generally have the same chemical properties (as reactivity is more influenced by protons/electrons). Generally, most atoms have a very similar number of neutrons and protons. For example, the common isotopes of carbon (Z=6) contain 6, 7, or 8 neutrons. The “default” configuration is an equal number of neutrons and protons, but this does not have to be the most common version in nature. The mass number (A) of an atom is the sum of its protons and neutrons. Electrons are not counted in this as their weight is so much less than protons and neutrons that it is insignificant. ATOMS AND ELEMENTS ATOMS AND ELEMENTS The number of neutrons can be found by subtracting the atomic number from the mass number: Neutrons = A – Z Isotopes are indicated by writing the mass number as a superscript before the chemical symbol, or by writing it after symbol with a hyphen. Carbon with A=13: 13C or C-13 This is pronounced “carbon thirteen” ATOMS AND ELEMENTS Hydrogen is the smallest element, with only one proton (Z=1). Hydrogen is also the only element that can exist without any neutrons. 1H has no neutrons: N=A–Z N=1–1 N=0 Hydrogen-1, commonly known as protium. ATOMS AND ELEMENTS Other isotopes of hydrogen have neutrons, however: H-2 (common name: deuterium) has 1 neutron: N=A–Z N=2–1 N=1 H-3 (common name: tritium) has 2 neutrons: N=A–Z N=3–1 N=2 Water molecules made up of H-2 rather than H-1 are called “heavy water” as the extra neutron per hydrogen atom increases the mass of the water. Heavy water has many industrial uses. RADIOACTIVITY Not all isotopes of an element are stable. Some isotopes are prone to decay into smaller pieces (either casting off some sub-atomic particles and/or splitting entirely into two or more nuclei). This is a nuclear reaction known as radioactive decay. A depiction of radioactive decay. In this image, an alpha particle (2 neutrons and 2 protons) is ejected from a larger nucleus. RADIOACTIVITY One useful type of radioactive decay is nuclear fission (fission = splitting). In this reaction, certain specific elements have their nuclei broken down into smaller parts. This reaction releases a tremendous amount of energy, which can be used for an explosion (nuclear weaponry), or to power and electric generator (nuclear reactor). U-235 is bombarded with a neutron to create U-236, which rapidly undergoes nuclear fission to yield Kr-92, Ba-141 and 3 lone neutrons. In many fission reactions, these by-products will themselves undergo fission until they reach a stable atomic isotope. RADIOACTIVITY Several isotopes undergo radioactive decay naturally at extremely steady rates and over very long periods of time. By knowing the original amount of a decaying isotope that was present and comparing it to the current amount, radiometric dating can be done to estimate the age of an item. Carbon-14 takes 60,000 years to fully decay. Carbon-14 dating allows scientists to date any item from within this time period. Naturally, items older than this cannot be C-14 dated. For older items, such as dinosaur fossils, other elements such as Uranium-238 are used, as their decay occurs over a much longer timeframe. U-238 dating is highly accurate, allowing us to measure things up to 4.5 billion years old to within 1% of error. RADIOACTIVITY Radiation exposure is very harmful to living things. The energy and small particles released from radioactive reactions will disrupt and destroy the biological molecules that make up living cells. This in turn will cause severe disruption of life processes, localized cell death, or complete death depending on the exposure. Radiation shielding uses materials (such as thick lead plates) that absorb the energy and particles released from radioactive processes. RADIOACTIVITY However, many radioactive isotopes have beneficial uses in medicine: Radioactive (but extremely harmless) technetium-99 can be injected into the body. As it circulates, radiation detectors can be used to observe its location and flow – helping to map out the body and detect circulation problems. Radioactively-tagged molecules (a otherwise-normal molecule containing a radioactive atom) are often used in research to find out what happens to molecules in living things. For example, a protein may be made containing radioactive sulfur and fed to an organism. Later, the scientists examine the organism and see what chemicals that organism has incorporated that sulfur into, thus tracking what substances that organism transforms its dietary protein into. RADIOACTIVITY Radiation therapy uses the destructive powers of radiation in as careful and specific a manner as possible. Radiation is used on extremely-specific regions of the body to kill cancer cells. Naturally, the radiation cannot differentiate between cancer and healthy cells and thus radiation therapy must be applied to as precise an area as possible to reduce the amount of healthy cells damaged. FUSION The opposite reaction to radioactive decay is that of nuclear fusion. Fusion, as the name implies, involves the combination of nuclei into a new, larger nucleus. Both nuclei are positivelycharged and thus will resist being pushed together. This makes the energy needed for fusion extremely high. Thermonuclear weapons create a small fusion reaction to unleash tremendous destructive power. Nuclear fusion of H-2 and H-3 into He-4 and a lone neutron. This reaction yields a large amount of energy (measured here in electronvolts (MeV)). FUSION After the big bang, the majority element was (and still is) hydrogen. Larger elements have been formed over time within stars. The immense size and pressure of a star’s core causes fusion to occur, creating larger and larger elements, depending on the size of the star. Larger elements are only found in regions of our universe that have had stars in them at some point. A depiction of the inner portion of a star showing subsequently heavier elements formed towards the center. Not to scale. PERCENT ABUNDANCE While many elements have multiple stable isotopes, that does not mean that all stable isotopes will exist in equal amounts in nature. The percent abundance is a measure of how much of a particular isotope exists compared to the whole of that element. The reasons for each isotope’s percent abundance vary from how much is made in solar furnaces, the formation of planets, planetary geology, and others. PERCENT ABUNDANCE On Earth, most elements will have at least two decently-stable isotopes that will both have fairly high abundance. This is why when we want to know the mass of an element, we do not use the mass of any one isotope – we use the average atomic mass, a weighted average of the various possible isotopes. For example, gallium exists as both Ga-69 (60.11%) and Ga-71 (39.89%). Therefore, if one wanted to know the mass of a piece of gallium, it would be wrong to assume the mass to be that of Ga-69 (exactly 69 amu) or Ga-71 (exactly 71 amu). The mass will be in between the two, though closer to 69, as it was the majority component. The average atomic mass of gallium works out to be 69.723 amu. PERCENT ABUNDANCE Average atomic mass is can be calculated using the formula: Where the fractional abundance is the percentage abundance of that isotope expressed as a decimal. This formula can be extended for as many isotopes are there are present. PERCENT ABUNDANCE PERCENT ABUNDANCE Example 2: Magnesium has 3 major isotopes: Mg-23 (23.99amu, 78.99%), Mg-24 (24.99 amu, 10.00%), and Mg-25 (25.98 amu, 11.01%). Find the Average Atomic Mass. AMM = (0.7899 x 23.99 amu) + (0.1000 x 24.99 amu) + (0.1101 x 25.98 amu) AMM = (18.9497 amu) + (2.499 amu) + (2.860398 amu) AMM = 24.31 amu. PERCENT ABUNDANCE The atomic mass given on the periodic table is the average atomic mass in amu (atomic mass units). Note that it is not a whole number for most elements as two or more isotopes exist. Many of the largest elements are listed with a wholenumber AAM, often in brackets, or have no AAM listed at all. These are elements that have only been created artificially, and thus have no “natural” abundance. The number listed is of the most-stable isotope created to date. IONS An element is defined by its number of protons. Neutrons may vary, but only contribute to the mass of the atom. In an atom, the number of electrons is always equal to the protons. This results in the positive and negative charges cancelling each other out, and thus atoms always have a net charge of zero. However, if the number of electrons and protons are different, the particle is called an ion. IONS Atom (charge = 0) Anion (charge < 0) Cation (charge > 0) IONS The net charge of an ion is the summation of the positive and negative charges (that is, the numbers of protons and electrons, respectively) in an ion. Net Charge = (protons) + (electrons) Note that electrons should be counted as negative in this formula. Net charge is usually shown to the right of a chemical symbol as a superscript. Ions are often shown encased in of square brackets. For atoms (which have no charge) this is naturally omitted. [F]- means a fluorine ion with a one minus charge. The “1” can be omitted or included. Other integers (2, 3, etc) must be indicated, such as in [O]2-. IONS To calculate the number of electrons in an ion, subtract the number of protons (the atomic number, Z) from the net charge. Electrons = Z – [net charge] For example, a O2- ion: 1) Oxygen has Z=8 (look up on periodic table) 2) A 2- ion has a net charge of -2. Electrons = (-2) – (8) Electrons = (-10) Therefore, there are 10 (negatively-charged) electrons in a O2- ion. SUMMARY Protons: Protons are determined by the atomic number, Z If Z is not given, find using the mass number (pro = mass – neu), or from net charge (pro = [net] – elec) Neutrons: Neutrons are found by using the mass number and the atomic number Neutrons = (mass number) – (protons) Electrons: Electrons are found by using the net charge and the number of protons Electrons = (net charge) – (protons) PRACTICE 1) How many electrons are in a Mg2+ ion? 2) What is the net charge on a Boron-10 ion/atom that contains 2 electrons? 3) What is the net charge on a Boron-10 ion/atom that contains 5 electrons? 4) What element contains 2 electrons and a net charge of 2+?