Unit 4: The Electromagnetic Spectrum, Energy & Electrons Notes

**Unit 4: The Electromagnetic Spectrum, Energy, & Electrons** **[Waves]** A wave is a [disturbance] that travels through space and matter, transferring [energy](https://www.ducksters.com/science/energy.php) from one place to another. It is important to remember that waves transfer energy, not matt...

**Unit 4: The Electromagnetic Spectrum, Energy, & Electrons** **[Waves]** A wave is a [disturbance] that travels through space and matter, transferring [energy](https://www.ducksters.com/science/energy.php) from one place to another. It is important to remember that waves transfer energy, not matter. There are many types of waves all around us in everyday life. - Sound is a type of wave that moves through matter and then vibrates our eardrums so we can hear. - You can drop a rock into a pond and see waves form in the water. - We even use waves (microwaves) to cook our food really fast. **[Types of Waves]** All waves can be categorized as either mechanical or electromagnetic: ***Mechanical waves*** are waves that require a [medium]. This means that they have to have some sort of matter to travel through. These waves travel when molecules in the medium collide with each other passing on energy. ***Electromagnetic waves*** are waves that are produced by the vibration of charged particles. These waves can travel through a [vacuum] (empty space); they don\'t need matter to travel through. Instead, they travel through electrical and magnetic fields that are generated by charged particles. \ Do you think that waves produced in the ocean are considered mechanical waves or electromagnetic waves? Why? [Ocean waves require the medium of water to travel through, so they are considered to be mechanical waves. ] Electromagnetic waves that are produced on the sun will travel to Earth through the vacuum of outer space. If this did not happen, there would be no life on Earth at all! Because these waves are so important to us and we see them everywhere, let's take a closer look at these waves:  **[Wavelength & Frequency]** Waves have two special properties: Wavelength and Frequency As shown in the above picture, wavelength and frequency have an [inverse] relationship: - The higher the **frequency**, the shorter the **wavelength** - The lower the frequency, the longer the wavelength Wavelength: - **Denoted by** λ **(Greek letter lambda**) - **Definition: the distance between crests (top of the wave) on adjacent waves** - Unit: meters, m **Frequency** - **Denoted by** ν **(Greek letter** nu) - Definition: The number of waves that pass a given point in 1 second - Units (all three units are the same): **Hertz (Hz) = 1/s = s^-1^** All electromagnetic waves travel at the speed of light (c = **3.00 x 10^8^ m/s). The equation that relates wavelength and frequency for electromagnetic waves is:** **λν = c OR λ = c/ν OR ν = c/λ** c = speed of light ; **λ =** wavelength; ν = frequency  Calculate the wavelength of a radio wave with a frequency of 1.027 x 10^8^ s^-1^ λ = c/ν = **3.00 x 10^8^ m/s ÷** 1.027 x 10^8^ s^-1^ = 2.921129503 m (3 sf) = 2.92 m Determine the frequency of light whose wavelength is 5.267 nm. 5.267 nm = 5.267 x 10^-9^ m (Use dimensional analysis) **[The Electromagnetic Waves]** **Radio:** Your radio captures radio waves emitted by radio stations, bringing your favorite tunes. Radio waves are also emitted by [stars](https://imagine.gsfc.nasa.gov/resources/dict_qz.html#star) and gases in space. **Microwave:** Microwave radiation will cook your popcorn in just a few minutes, but is also used by [astronomers](https://imagine.gsfc.nasa.gov/resources/dict_ad.html#astronomy) to learn about the structure of nearby [galaxies](https://imagine.gsfc.nasa.gov/resources/dict_ei.html#galaxy). **Infrared:** Night vision goggles pick up the infrared light emitted by our skin and objects with heat. In space, infrared light helps us map the [dust](https://imagine.gsfc.nasa.gov/resources/dict_ad.html#dust) between stars. **Visible:** Our eyes detect visible [light](https://imagine.gsfc.nasa.gov/resources/dict_jp.html#light). Fireflies, light bulbs, and stars all emit visible light. **Ultraviolet:** Ultraviolet radiation is emitted by the Sun and are the reason skin tans and burns. \"Hot\" objects in space emit UV radiation as well. **X-ray:** A dentist uses x-rays to image your teeth, and airport security uses them to see through your bag. "Hot" gases in the [Universe](https://imagine.gsfc.nasa.gov/resources/dict_qz.html#universe) also emit X-rays. **Gamma ray:** Doctors use gamma-ray imaging to see inside your body. The biggest gamma-ray generator of all is the Universe. **[Introduction to Light ]** The sun is constantly producing energy that reaches the Earth's surface. The majority of energy that we receive from the sun is in the form of heat and what we call "light." Light usually refers to the way your eyes perceive things. They can detect when light hits an object, or they can detect when an object emits (gives off) light. The common use of light refers to electromagnetic waves that are **visible** to the naked human eye (the visible light spectrum). The other six types of electromagnetic waves cannot be seen with the naked human eye, but these waves are considered to be light too. The interesting thing about light is that it behaves like both [waves] (with frequencies and wavelengths) and [particles] (like the subatomic particles in the atom). We have already discussed waves, so let's look at particles next. **[Quantum Theory & The Photoelectric Effect]** Max Planck (early 1900s) stated that light can be emitted by objects when the energy of that object was changed by small increments called [quanta]. A [quantum] of energy is the smallest amount of energy that can be lost or gained by an atom. These observations are called the Quantum Theory. Planck's work was further supported by Albert Einstein. Einstein observed that if high frequency light (gamma rays) were directed towards the surface of a metal, these waves would carry enough energy to the electrons in the metal, causing these electrons to vibrate, and then the electrons would be [ejected] (thrown out) from the metal. Einstein discovered that the energy from light is a particle (photon), and that these photons traveled in a wave motion. Therefore, **light is both a wave and a particle**! This observation is called The Photoelectric Effect.  **[Radiation (Energy)]** How are the electromagnetic waves related to each other? They are produced and detected in different ways, but all waves on the electromagnetic spectrum produce light and radiation. Electromagnetic radiation is produced by electromagnetic waves that are emitting energy. What is emitting this energy? It is the [electrons], as described in The Photoelectric Effect. Electrons vibrate and emit energy in the form of [photons](https://imagine.gsfc.nasa.gov/resources/dict_jp.html#photon). Photons, like electrons, have no mass and they are pure energy. The electromagnetic waves are made up of photons. Photons travel in a wave-like pattern at the [speed of light](https://imagine.gsfc.nasa.gov/resources/dict_qz.html#speed_of_light) (c = **3.00 x 10^8^ m/s**) and carry a certain amount of energy, which results in the different types of radiation. **To calculate the energy carried by an electromagnetic wave, use the formula: E = hv or E = hc/ λ** E = Energy (Units: Joules, J) h = Planck's constant = 6.626 x 10^-34^ Js (Units: Joules x second) v = Frequency (Units: **Hz, 1/s, s^-1^)** What is the energy of green light with a wavelength of 500 nm? Convert 500 nm to 500 x 10^-9^ m (using dimensional analysis). Then use the formula E = hc/λ since you are given a wavelength E = (6.626 x 10^-34^ Js)(3.00 x 10^8^ m/s) **÷ (**500 x 10^-9^ m) = 4 x 10^-19^ J (1 sig fig) **[Energy, Frequency, & Wavelength]** Energy has a [direct] relationship to frequency, so it will also have an [inverse] relationship to wavelength. Electromagnetic radiation can be expressed in terms of energy, wavelength, or [frequency](https://imagine.gsfc.nasa.gov/resources/dict_ei.html#frequency).  **[Bohr Models Revisited]** Niels Bohr modified the previous atomic model made by Ernest Rutherford by stating at electrons move in circular orbits around the nucleus. Bohr based his model off of research done on the [Hydrogen atom]. The rings in the Bohr Model are called energy levels. Bohr used his model to show how the energy of an atom changed when the atom absorbed or emitted light. [Absorption] occurs when the photons (energy) of light are taken in by electrons, causing the electrons to become "excited" and move into higher energy levels. On the Bohr Model, the electrons would move away from the nucleus as the electrons get excited. [Emission] occurs when the photons are released by the electrons, causing the electrons to fall back into the "ground" state. On the Bohr Model, the electrons would move towards the nucleus as the electrons lose energy. The energies absorbed and emitted are exactly the same. **[Visible Light Spectrum]** Visible light corresponds to light with wavelengths ranging from around 400 nm to 700 nm (it is easier to use nanometers than meters here). The Sun produces lights with all of these wavelengths mixed together, which appears to our eyes as [white light]. When this white light passes through a [prism], the light separates into the individual colors of the visible light spectrum:  Violet color has wavelengths of 400 -- 430 nm Indigo color has wavelengths of 430 -- 460 nm Blue color has wavelengths of 460 -- 500 nm Green color has wavelengths of 500 -- 560 nm Yellow color has wavelengths of 560 -- 595 nm Orange color has wavelengths of 595 -- 620 nm Red color has wavelengths of 620 -- 700 nm What can you say about the frequency and energy of red light and violet light based off the above wavelengths? [Red light has the longest wavelength, but will have the lowest frequency and energy. Violet light has the shortest wavelength, but will have the highest frequency and energy.] **[Atomic Emission Spectrum]** Radiation is the energy produced when an atom absorbs energy, causing the electrons to move into higher energy levels (ground state → excited state). When these electrons emit this energy (excited state → ground state), an electromagnetic wave is produced. Depending on the type of atom and the amount of energy, the radiation emitted can take the form of [visible light], one of the other electromagnetic waves, or heat. Think about neon signs. How are we able to see this light? [Neon (a gas) is contained in tubes. Electricity applied to the neon caused the electrons to become excited. After neon's electrons lose this energy, we see the loss in the form of light.] As seen in the above neon example, one way for atoms to absorb/emit light is by applying an [electric current] to the [gaseous] state of an element. The light that is emitted after the atoms begin to lose energy can be seen as light. Like with sunlight, a prism can be used to separate this light into a [line spectrum], which is also called an atomic emission spectrum. Every element has a unique line spectrum! Shown below are the line spectra for five elements. The last line spectrum is from an alloy containing one or more of these five elements. Identify which elements are present in the alloy.  [By lining up the individual lines in the unknown spectra with lines on each of the elemental line spectra, it seems that the alloy contains cadmium, sodium, and strontium, Lithium and calcium are ] [not present. ] Two atomic models are in use today: the Bohr Model and Schrödinger's Quantum Mechanical Model. The Bohr Model can be used to understand vital concepts such as absorption/emission. However, the Quantum Mechanical Model is a truer representation of the atom. **[Quantum Mechanical Model]** The Quantum Mechanical Model is related to The Quantum Theory proposed by Max Planck, which states that atoms absorb and emit energy in specific amounts called quanta. Einstein further stated that it is the electrons that absorb and emit quanta. However, it is impossible to determine both the [location] and the [speed] of the electron at the same time. This is called the [Heisenberg Uncertainty Principle]. The Quantum Mechanical Model is a mathematical equation. Solutions to this equation show the approximate location of an electron. This location is one of many spaces, called [orbitals], that the electron can be found in. All of the orbitals combined together make up the [electron cloud]. **[Quantum Numbers]** The Quantum Mechanical Model generates four numbers that describe electrons and their orbitals: - - Principle Quantum Number: n - Angular Momentum Quantum Number: ℓ - Magnetic Quantum Number: m~ℓ~ - Spin Quantum Number: m~s~ Principle Quantum Number (n): This number describes the distance of an orbital from the nucleus. It is also known as the [energy level] of the electron. The possible values for n are 1, 2, 3, 4, 5, 6, or 7. The energy level corresponds to the period (on the periodic table) or the energy level (from the Bohr Model). Angular Momentum Quantum Number (ℓ): This number describes the different energy [sublevels] of the orbital and the orbital shapes that the electron can travel on. The four orbitals are s, p, d, and f. Each of the orbitals has a different number of orbital shapes. Each orbital can hold electrons. - The "s" block has 1 possible orbital shape, so it can hold electrons. - The "p" block has 3 possible orbital shapes, so it can hold electrons. - The "d" block has 5 possible orbital shapes, so it can hold electrons. - The "f" block has 7 possible orbital4shapes, so it can hold electrons. **Refer to the attached handout for the "n" and "s, p, d, f" regions**  To remember how many electrons each orbital can hold, count the number of elements in that region! Draw the Bohr Model for Potassium. Quantum studies have led to three rules that govern the way we describe electrons: 1. Aufbau Principle - [Electrons must occupy the lowest energy level first before moving into the next higher energy level] - This is similar to how we filled in the energy levels in the Bohr Model (fill in level one and then move into level two) 2. Hund's Rule - [Electrons must fill each orbital before it can pair up with another electron.] - This is similar to how we filled in the energy levels in the Bohr Model in rings 2, 3, and 4 above (draw an electron on each of the four corners before going back and pairing up electrons). 3. Pauli Exclusion Principle - [Only two electrons can occupy a single orbital. The first electron has an upwards spin, and the second electron has a downwards spin.] **[Electron Configuration]** Electron configurations tell us all the orbitals the electrons for an element occupy. To write the electron configuration for an element, you must take into account all of the above rules AND all of the elements that come before that element on the periodic table. You can use the following tree diagram to guide you: n**s**^e^ n**p**^e^ n**d**^e^ n**f**^e^ n will be the energy level (period) - s and p: n = period - d: n = period -- 1 - f: n = period - 2 e will be the number of electrons found in that region You always start writing electron configurations with Hydrogen, and then continue from there until you reach your element. Write the electron configurations for the first five elements and titanium. What do you notice about all of the exponents when you add them together? H: 1s^1^ [Adding all the exponents together gives the atomic number!] He: 1s^2^ Li: 1s^2^2s^1^ Be: 1s^2^2s^2^ B: 1s^2^2s^2^2p^1^ Ti: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^2^  Write the electron configurations for elements 11, 12, 15, 20, 26, 34, 45, and 83. Na: 1s^2^2s^2^2p^6^3s^1^ Mg: 1s^2^2s^2^2p^6^3s^2^ P: 1s^2^2s^2^2p^6^3s^2^3p^3^ Ca: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^ Fe: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^6^ Se: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^4^ Rh: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^6^5s^2^4d^7^ Bi: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^6^5s^2^4d^10^5p^6^6s^2^4f^14^5d^10^6p^3^ **[Noble Gas Electron Configuration]** Noble Gas Electron Configuration is a shorthand way of writing electron configurations for elements (as you saw with the previous examples, these configurations can become really long). To use Noble Gas configuration: - Locate the noble gas that comes before the element you are writing the configuration for - Write the noble gas in the following format: \[Symbol\] - Continue writing the configuration for everything that follows that Noble Gas Write the electron configuration for Titanium (from page 10). Then, write the Noble Gas electron configuration for Titanium.  Write the Noble Gas configurations for elements 11, 12, 15, 20, 26, 34, 45, and 83. Na: \[Ne\]3s^1^ Mg: \[Ne\]3s^2^ P: \[Ne\]3s^2^3p^3^ Ca: \[Ar\]4s^2^ Fe: \[Ar\]4s^2^3d^6^ Se: \[Ar\]4s^2^3d^10^4p^4^ Rh: \[Kr\]5s^2^4d^7^ Bi: \[Xe\]6s^2^4f^14^5d^10^6p^3^ **[Orbital Diagrams]** Orbital Diagrams are similar to electron configurations, except you can see how the electrons are distributed within each of the subshells (s, p, d, f). 1. Start by writing the electron configuration of the element. 2. Draw lines for the number of orbitals for each subshell (s has 1, p has 3, d has 5, f has 7). 3. Label the lines by writing the subshell beneath the line. 4. Fill in electrons using the three rules from page 9. Draw the orbital diagram for Titanium. Ti: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^2^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s 3d Draw the orbital diagram for elements 11, 12, 15, 20, 26, 34, 45, and 83. Na: 1s^2^2s^2^2p^6^3s^1^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s Mg: 1s^2^2s^2^2p^6^3s^2^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s P: 1s^2^2s^2^2p^6^3s^2^3p^3^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p Ca: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s Fe: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^6^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s 3d Se: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^4^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s 3d 4p Rh: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^6^5s^2^4d^7^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s 3d 4p \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 5s 4d Bi: 1s^2^2s^2^2p^6^3s^2^3p^6^4s^2^3d^10^4p^6^5s^2^4d^10^5p^6^6s^2^4f^14^5d^10^6p^3^ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 1s 2s 2p 3s 3p 4s 3d 4p 5s \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ \_\_ 4d 5p 6s 4f 5d 6p