Name: rusting and metal protection
Uploaded: 2020-09-24T15:06:52.000Z
Duration: 2 min 49 s
Channel: JoVE

K12 Standards

Arizona Science Standards

High School

Physical Science

HS.P1

HS.P1U3.4

Corrosion

腐蚀是金属在氧化剂存在下的 自发氧化，氧化剂一般是指氧气。自由女神像，或生锈的锚，会出现蓝绿色，也是这种现象。有些腐蚀具有保护作用。例如，铜在 氧气、二氧化碳和水 的存在下氧化，形成碱式碳酸铜，这是一种 蓝绿色且结构稳定的保护层。其他腐蚀是有害的。例如，铁在氧、水 和氢离子的存在下氧化成 结构不稳定的金属氧化物，称为铁锈，它会不断剥落，使下面的金属进一步氧化。生锈是一种自发进行的电化学反应。表面缺陷成为铁的阳极区，电极电位为 0.44 伏，很容易氧化为二价铁。电子向阴极区迁移，阴极区的电极电位为 1.23 伏，氧气在酸性环境中被还原成水。氢离子来自大气中二氧化碳和水的反应，生成碳酸。接下来，二价铁离子通过表面水分迁移至阴极区，进一步氧化为三价铁氧化水合物，通常称为"铁锈"水合物的化学计量随二价铁 暴露的水量而变化。多余的水分、酸和电解质 会加速生锈，从而增加电荷流动 和 H+离子的浓度。防止腐蚀的方法包括 在金属表面涂漆，防止 与水和氧气接触，或使用 由多种金属组成的合金。不锈钢一种含有少量铬的铁合金，经过腐蚀后会在铁表面形成一层 不活泼的氧化层来保护铁。另外，还可以采用电镀方法。这里，铁螺丝上涂有一层 更容易氧化的金属，例如锌，形成一层氧化锌保护层 来保护铁。最后，还有阴极保护法，它通过把被保护的金属转变为阴极来控制腐蚀。地下铁管通常连接到更容易氧化的金属上，例如连接到锌上，锌在这里充当被氧化的牺牲阳极，而铁则成为免受腐蚀的阴极。

腐蚀

自然电化学过程导致金属降解称为腐蚀。 铁锈形成，银渍蚀以及铜上出现的蓝绿色专利就是腐蚀的例子。 腐蚀涉及金属氧化。 有时它是保护性的，例如铜或铝的氧化，其中金属氧化物或其衍生物的保护层形成在表面，从而保护底层金属不会进一步氧化。 在其他情况下，腐蚀会损坏金属，例如生锈铁。 
 不良氧化还原反应：铁锈 
 由于铁暴露在氧气和水中，生锈。 铁锈形成包括在铁表面形成原电池（伽凡尼电池），从而产生铁 (II)。 在铁表面上形成的厌氧 (铁氧化) 和阴极 (氧减少) 区域发生的相关氧化还原反应包括： 
 <img alt="Eq1" src="/files/ftp_upload/11440/11440_Equation_1.png" style="text-align: center; height: 150px;" />
 铁 (II) 进一步与潮湿空气发生反应，形成一种铁 (III) 氧化水合物，通常称为锈。 
 <img alt="Eq2" src="/files/ftp_upload/11440/11440_Equation_2.png" style="text-align: center; height: 50px;" />
 水合物的化学物质化学计量因铁 (II) 暴露的水的数量而异，化合物公式中使用 x 表示。 潮湿，酸和电解质的存在会增加生锈形成的率值。 与铜上的铁杆不同，铁锈的形成不会形成保护层，因此随着铁锈的脱落，新鲜铁暴露在大气中，铁的腐蚀将继续存在。 
 防腐蚀 
 可以使用各种方法来防止腐蚀。 一种方法是保持金属表面的油漆，以避免与水和氧气接触。 金属合金，如将铁与少量铬混合在不锈钢中，是防止腐蚀的另一种有效方法。 铬在表面附近聚集，并被氧化，从而有效地保护铁不受腐蚀。 
 铁和其他金属也可以通过电镀来防止腐蚀，在过程中，要保护的金属涂有一层更容易氧化的金属，通常是锌。 当锌层完好无损时，它可以防止空气接触底层铁，从而防止腐蚀。 如果锌层被腐蚀或机械磨损破坏，则仍可通过阴极保护过程来保护熨斗免受腐蚀，下一段对此进行了描述。 
 阴极保护利用将要保护的金属转换为电化学反应中的阴极的原理。 通过将受保护金属连接到更有效或更容易氧化的金属 (如锌或镁) (称为牺牲阳极) 来实现这一点。 阳极腐蚀并被用于保护作为阴极的金属。 阴极保护最常用于家用电器，如热水器和地下储水箱。 重要的是，阴极保护可用于除铁以外的金属。 
 本文改编自<a href="https://openstax.org/books/chemistry-2e/pages/17-6-corrosion"> OpenStax, 化学 2e, 第 17.6 章：腐蚀</a>。

由于自然电化学过程导致的金属降解被称为腐蚀。铁上形成的铁锈、银的失去光泽以及蓝绿色的……

Education

教育

JoVE Core

Chemistry

化学

Electrochemistry

电化学

KEY TERMS AND CONCEPTS

关键术语与概念

rusting and metal protection

Corrosion is the natural breakdown of metals through electrochemical reactions. It can appear as rust on iron, tarnish on silver, or a blue-green patina on copper. In each case, the metal is oxidized. Some corrosion products can protect the surface, while others keep damaging the metal.

Rusting is the harmful type of corrosion seen in iron. It happens when iron is exposed to oxygen and water. A galvanic cell forms on the iron surface, and this drives oxidation at the anodic region and reduction of oxygen at the cathodic region. The iron(II) that forms then reacts with humid air to make iron(III) oxide hydrate, which is rust.

The amount of water in rust can vary, so its formula includes x to show changing stoichiometry. Moisture, acids, and electrolytes all speed up rust formation. Unlike the protective patina on copper, rust does not seal the metal. It flakes away and exposes fresh iron to air, so corrosion continues.

Metal corrosion can be slowed in several ways. One method is painting the surface so water and oxygen cannot reach it. Another is alloying, such as adding chromium to iron to make stainless steel. The chromium moves near the surface and oxidizes there, which helps protect the iron underneath.

Metals can also be protected by galvanization, where a layer of a more easily oxidized metal, usually zinc, coats the surface. If the zinc layer stays intact, it blocks air from the iron. If the coating is damaged, cathodic protection can still help. In that method, the protected metal is connected to a more active sacrificial anode, such as zinc or magnesium, which corrodes instead. This approach is used in water heaters and underground water storage tanks, and it can protect metals other than iron.

Rusting and Metal Protection

balancing redox equations

Redox Reactions and Half-Reaction Balancing

Redox reactions involve electron transfer and changes in oxidation number. In electrochemistry, these reactions connect chemical changes with electrical energy. They also power technologies such as industrial metal processing and rechargeable batteries used in electric vehicles. Oxidation number is a way to track how an atom's electronic environment changes compared with the pure element, where the oxidation number is zero. For a simple ionic compound, the oxidation number matches the ion...

electromotive force

Cell Voltage and Electron Flow

Electric current is the flow of electric charge through a conducting medium or a solution. In wires, that charge is carried by electrons. In some solutions, ions can also move and carry charge. When electrons move through a wire, they produce an electric current. This idea connects to redox reactions, where electrons are lost and gained. In the spontaneous reaction between zinc and copper ions, zinc is oxidized to zinc ions. Copper ions are reduced to solid copper. Even so, this direct...

voltaic/galvanic cells

Galvanic Cells and Battery Reactions

Galvanic cells use spontaneous redox reactions to produce electric current. A common example is a disposable AA battery, which powers devices by turning chemical energy into electrical energy. Another example is a copper wire placed in silver nitrate solution, where the solution slowly turns bright blue and a grey solid forms on the wire. In the copper and silver reaction, copper is oxidized and silver ions are reduced. Copper forms cuprous ions, which give the solution its blue color. At the...

standard electrode potentials

Using Electrode Potentials to Predict Reactions

Electrode potentials help predict whether a redox reaction will happen on its own. A redox reaction involves electron transfer. In electrochemistry, the measured value is called cell potential or cell voltage. It shows the energy linked to charge transfer and is measured in volts, where 1 volt equals 1 joule per coulomb. A single half-cell cannot be measured by itself because electron transfer needs both a donor and a receiver. For that reason, scientists measure the difference between two...

cell potential and free energy

Redox Cell Potential, ΔG°, and K

Redox cell potential connects chemical energy, free energy change, and the equilibrium constant in an electrochemical cell. In these reactions, chemical energy is converted into electrical energy. That link helps predict whether a reaction is spontaneous and how far it may proceed. Thermodynamics is the study of how heat and other forms of energy are related. In an electrochemical cell, cell potential can also reflect oxidizing strength or reducing strength. A similar idea appears in acid-base...

the nernst equation

Cell Potential Under Changing Conditions

Cell potential changes when a redox reaction does not run at standard conditions. Standard cell potentials connect with standard free energy change, ΔG°, and the equilibrium constant K, but those relationships only hold at 1 M concentration, 298 K, and 1 atmosphere. A zinc(II) and tin(II) redox cell with E°cell = +0.291 V and ΔG° = −56.2 kJ is spontaneous under those standard conditions. As a cell discharges, the reactant concentration changes and the cell potential steadily drops. Many redox...

concentration cells

Cell Potential from Concentration Differences

Concentration cells are voltaic cells that use two almost identical half-cells. The half-cells use the same electrode and the same half-reaction. They differ only in the concentration of one redox species, so the cell potential depends on that concentration difference. In a concentration cell, the Nernst equation can be used to find the cell potential. For the zinc example shown, the calculated cell potential is positive. A positive cell potential means the overall reaction is spontaneous. <br...

batteries and fuel cells

Primary, Rechargeable Batteries and Fuel Cells

Batteries and fuel cells are galvanic cells that supply electrical power for many applications. They range from tiny button batteries in wristwatches to large systems that can provide backup energy to municipal power grids. Some are primary cells, which are single-use and cannot be recharged. Others are secondary cells, which use reversible reactions and can be charged again by an external power source. A common primary battery is the dry cell. It uses zinc as both the container and the anode,...

corrosion

Corrosion is the natural breakdown of metals through electrochemical reactions. It can appear as rust on iron, tarnish on silver, or a blue-green patina on copper. In each case, the metal is oxidized. Some corrosion products can protect the surface, while others keep damaging the metal. Rusting is the harmful type of corrosion seen in iron. It happens when iron is exposed to oxygen and water. A galvanic cell forms on the iron surface, and this drives oxidation at the anodic region and...

electrolysis

Electrolysis in Batteries and Metal Production

Electrolysis uses electrical energy to drive a nonspontaneous redox reaction. In a galvanic cell, a spontaneous redox reaction does electrical work on the surroundings by sending electrons through an external circuit. In electrolysis, an external circuit does work on the redox system by applying a voltage large enough to force the reaction in the opposite direction. This process is used in several real-world settings. Recharging a battery is one example, because an outside power source drives...

key terms and concepts

Learn how corrosion affects metals, why iron rusts, and how painting, stainless steel, galvanization, and cathodic protection help prevent damage.

Rusting and Metal Protection | JoVE K12

chemistry

JoVE Core 化学通过简单明了且通俗易懂的动画视频课程，将普通化学概念带入到现实生活当中。此外，在科学家行动视频中，我们演示了在实验室中进行的经典和原始研究实验。

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Rusting and Metal Protection

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